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汉语一级 > chemistrah > Flashcards

chemistrah Flashcards

1
Q

give l values for s, p, d, and f subshells

A 
l=
0
1
2
3
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2
Q

max number of electrons for a single orbital

A

2

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3
Q

describe degenerate orbital

A

same energy level but pairing takes more energy so the suborbitals will be formed with symmetrically unpaired electrons before they pair up

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4
Q

paired electrons have ______ spin

A

opposite

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5
Q

determine molecular formula given empirical formula –

A

find approximate formula mass (add up the masses of component atoms)

if you know the total molecular mass, divide by the formula mass

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6
Q

how many total atoms are there in 53.9g hydrazine?

A

hydrazine = n2h2

number of molecules in hydrazine:
1.08 x 10^24

this times four (four atoms per molecule) should give 4.32x10^24

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7
Q

if an ion says + it has _____ an electron

A

lost

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8
Q

if an ion says - it has _____ an electron

A

gained

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9
Q

how many electrons can the third principal energy level hold?

A

18

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10
Q

define resonance

A

multiple equivalent or nearly equivalent Lewis structures can be drawn for a molecule

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11
Q

what is the reactivity of compounds with resonance (compared to what is expected)

why?

A

compounds with resonance are generally less reactive than would be expected from the number of multiple bonds

the spreading of a π bond is called de-localisation

compared to localised π electrons, which are shared between only two atoms, delocalised π electrons are much less available for reactions

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12
Q

bond order =

A

1/2 (bonding e- – antibonding e-)

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13
Q

bond orders greater than 0 indicate

A

a stable particle

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14
Q

define paramagnetism

A

weak attraction to magnetic field

property of molecules with unpaired electrons

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15
Q

[MO theory] paramagnetism is indicated by ?

A

paramagnetism is indicated by the presence of one or more unpaired electrons

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16
Q

define diamagnetism

A

very slight repulsion from magnetic field

property of molecules with all paired electrons

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17
Q

diamagnetism results when

A

all electrons are paired

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18
Q

atoms hybridise when

A

bonding is imminent AND when the energy cost of forming the hybrid orbitals is offset by lower energy of final molecule

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19
Q

sp3 hybrid orbitals are formed by _____ and have

____ e- geo

A

one s and 3 p orbitals

tetrahedral

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20
Q

take carbon e.g.

energies of four resulting sp3 hybrid orbitals are _____ that of 2s and 2p atomic orbitals

in carbon how many electrons does each hybrid orbital contain?

A

BETWEEN

one, unpaired

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21
Q

sigma bonds

A

head-to-head overlap along internuclear axis

single bonds

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22
Q

double bonds are made up of (in terms of σ and π)

A

1 σ 1 π

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23
Q

sp hybrid orbitals are formed by _____ and have

____ e- geo

A

one s and one p orbital

linear

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24
Q

π bonds ?

A

side on overlap of unhybridised p orbitals above and below internuclear axis

found in double and triple bonds

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25
Q

sp2 hybrid orbitals are formed by _____ and have

____ e- geo

A

one s two p

trigonal planar

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26
Q

triple bonds are made up of (in terms of σ and π)

A

1 σ and 2 π

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27
Q

give e- geos

sp
sp2
sp3

A

linear
trigonal planar
tetrahedral

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28
Q

do you ever find d hybridized orbitals?

why or why not

A

not really

within a major energy level, the d sublevel energy is far higher than the s and p sublevel energies

so these combinations are energetically unfavourable

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29
Q

hybrid orbitals, sp3 sp2 etc, depend on which, electron or molecular geometries?

A

ELECTRON GEOMETRIES

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30
Q

Describe the relationship between bond length and bond strength

(Explain)

A

Bond length and bond strength are inversely related.

As bond enthalpy increases, bond length decreases; so stronger bonds hold atoms more tightly together.

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31
Q

enthalpy is

A

enthalpy is a stand-in for energy in chemical systems;

bond, lattice, solvation and other “energies” in chemistry are actually enthalpy differences.

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32
Q

In the International System of Units (SI), the unit of measurement for enthalpy is

others used include

A

the joule

the calorie
BTU british thermal units

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33
Q

Enthalpies of chemical substances are usually listed for _________ pressure as a standard state.

A

1 bar (100 kPa)

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34
Q

Rank the following bonds between H−C, H−O, and H−N from longest to shortest.

What you need to do here is explain correctly the reason for the order

A

H−C, H−N, H−O

Moving across the periodic table from left to right, atomic radius decreases. The smaller the atomic radius, the closer the atom will be to hydrogen and the shorter the bond.

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35
Q

The formal charge of an atom is determined with this formula:

A

formal charge = (# of valence electrons) − (# of nonbonded electrons) − 1/2(# of bonding electrons).

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36
Q

Which kind of bond exists between atoms of silicon and chlorine, Si−Cl?

A

polar covalent

Polar covalent bonds occur between two atoms with a large electronegativity difference. The electronegativities of silicon and chlorine are significantly different, so there is unequal sharing of the bonding electrons.

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37
Q

Explain the relationship between atomic size and bond strength.

A

The larger two atoms are, the further apart they must sit, and therefore the longer their bond will be.

Longer bonds are weaker bonds

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38
Q

Which kind of bond exists between atoms of carbon and hydrogen, C−H?

A

A C−H bond is covalent. Covalent bonds occur between two nonmetals with similar electronegativities. Electrons are shared equally in a pure covalent bond.

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39
Q

molecular compounds contain only

A

nonmetals

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40
Q

give hybrid orbitals for each of the following ELECTRON geometries

linear
trigonal planar
tetrahedral

A

sp
sp2
sp3

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41
Q

sp
sp2
sp3

describe electron geometries

A

linear
trigonal planar
tetrahedral

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42
Q

atoms located ____________ on the periodic table (such as ______) are capable of having an expanded octet

A

below the second period

sulphur e.g.

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43
Q

___________ can potentially have more than 8 electrons in valence shell

A

atoms located below the second period on periodic table

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44
Q

the larger the ___________, the greater the dipole moment

A

EN difference

(dipole moment depends on:
the length of the bond
**the electronegativities of the atoms

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45
Q

dipole moment – can it be measured?

A

yes it is measurable

units of debye (D) – includes the charge unit, coulomb, and the distance unit, metre

indicates how much ionic character is present in covalent bond

PERCENT IONIC CHARACTER

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46
Q

percent ionic character !!!!!!!!

A

compare measured dipole moment (μ measured)
to
calculated dipole moment (μ calculated) for 100% electron transfer (indicating an ionic bond)

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47
Q

charge of electron

A

1.6 x 10^-19 C

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48
Q

μ calculated

A

bond length and charge of electron (1.6 x 10^-19 C)

gives dipolemoment as if electron were completely transferred –> indicating 100% ionic character

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49
Q

formal charge =

A

formal charge = #val e- – (1/2)(shared e-) – #unshared e-

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50
Q

define hypervalence

A

has an expanded octet

below period 2

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51
Q

compound that contains carbon and hydrogen (and not much else) is called a

A

hydrocarbon

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52
Q

dry ice is __________

when heated it will ________

A

solid CO2

sublime

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53
Q

fire extinguishers spray out

how does it work

A

CO2
heavier than O2
displace O2
suffocate flame

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54
Q

decomposition reaction

A

something complex becomes simpler parts

AB –> A + B

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55
Q

single replacement (just use ABC)

A

A + BC –> AC + B

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56
Q

two types of double displacement reactions

A

acid-base reaction

precip reaction

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57
Q

combustion reaction basics

A

CxHy + O2 –> CO2 + H2O

hydrocarbon exposed to oxygen, quantities of carbon dioxide and water

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58
Q

combination reaction aka

A

synthesis reaction

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59
Q

enthalpy

A

heat energy that’s transferred

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60
Q

what makes a reaction occur?

A

both enthalpy and entropy

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61
Q

spontaneous reactions have

A

a favourable enthalpy and entropy combination

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62
Q

there are patterns that show a driving force in some reactions

A

these form patters that are energetically more favourable

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63
Q

acid base reaction aka

A

neutralisation reaction!

salt and water are energetically more stable than acid and base. this potential stability drives reaction towards products

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64
Q

single replacement driving force?

A

oxidation-reduction is the driving force

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65
Q

precipitation reaction is an example of

A

double displacement where precipitation is the driving force

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66
Q

decomposition reaction driving force?

A

oxidation reduction

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67
Q

soluble means

A

readily dissolves in water

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68
Q

insoluble means

A

doesn’t dissolve

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69
Q

denote soluble in formulae

A

formula (aq)

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70
Q

electrolytes are substances that

A

when dissolved in water conduct electricity

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71
Q

ionic compounds are known as

A

strong electrolytes bc they dissociate 100% in water

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72
Q

nonelectrolytes dissolve as

A

molecules not ions

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73
Q

acids are

A

molecular compounds that ionise

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74
Q

strong acids ionise ____

A

100%

dissociate completely making a strong electrolyte

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75
Q

list the 7 acids that dissociate completely

A 
halogens:
hydrochloric
hydrobromic
hydroiodic
&
perchloric
chloric

nitric acid
sulphuric acid

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76
Q

one of the strong acids has 2 hydrogens

A

h2so4

sulphuric acid

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77
Q

list 7 strong acids’ chem formulae

A 
HCl
HBr
HI
HClO4
HClO3

HNO3
H2SO4

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78
Q

list ions the 7 strong acids make in solutions

A 
H+  Cl-
H+  Br-
H+  I-
H+  ClO4-
H+  ClO3-

H+ NO3-

H+ HSO4-

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79
Q

an 8th strong acid?

A

chlorous

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80
Q

weak acids, define

A

dissociate only partially in water

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81
Q

what percent of weak acids ionise

A

only about 1%

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82
Q

strong bases, define

A

ionic compounds containing hydroxide (OH-) ions

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83
Q

what happens to strong bases in water

A

dissociate 100%

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84
Q

are strong bases electrolytes?

A

yes, they are strong electrolytes

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85
Q

list a couple of strong bases

A

NaOH, Ca(OH)2, KOH, Ba(OH)2

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86
Q

diff between strong and weak bases (3 things)

A

strong bases – ionic compounds
weak bases – molecular compounds

strong bases – dissociate 100%
weak bases – react with water to a small extent to produce OH ions (about 99% remaining as eg water and molecular ammonia)

strong bases – strong electrolytes
weak bases – weak electrolytes

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87
Q

do insoluble compounds dissolve?

A

yes, to a very small degree. ignore it for now

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88
Q

rules for assigning oxidation numbers – give rules 4 through 8

A 
4. grp 1 = +1, grp 2 = +2, Al = +3
F = -1
H = +1
O = -2
8. other halogens = -1, assign most EN first
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89
Q

give first three rules for assigning oxidation numbers

A
  1. neutral element, ON = 0
  2. monatomic ion, ON = ionic charge
  3. formula = 0 = sum ON within
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90
Q

review once more: assigning oxidation numbers rules 4 thru 8; only need to give affected elements (ONs are intuitive the order is not)

A 
grp 1, grp 2, Al
F
H
O
halogens >EN
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91
Q

which of the following are redox reactions?

double displacement
single displacement
synthesis

A

double displacement – never
single displacement – always
synthesis – some (often)

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92
Q

define combustion

A

rapid combination of a substance with oxygen

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93
Q

acid base reaction describe

A

acid + base –> salt + water

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94
Q

driving force in acid-base reaction?

A

driving force: neutralisation

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95
Q

driving force in precipitation reaction?

A

driving force: formation of a precipitate

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96
Q

A reaction involving a transfer of electrons from a higher energy state to a lower energy state; a driving force in many types of reactions:

A

oxidation-reduction

aka redox

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97
Q

redox is the driving force in three types of reactions

A

Single-replacement, synthesis, and decomposition

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98
Q

describe redox as a driving force

A

the transfer of electrons to form lower energy products

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99
Q

Reactions occur spontaneously due to a combination of

A

changes in heat energy (enthalpy) and randomness (entropy)

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100
Q

list the strong acids

A

7
HCl
HBr
HI

HClO4
(HClO3)

HNO3
H2SO4

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101
Q

molarity equation

A

M = moles / litre of solution

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102
Q

alternate molarity equation

A

M = millimole (mmol) / millilitre

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103
Q

list some strong bases

A 
LiOH
NaOH
KOH
Ca(OH)2
Sr(OH)2
Ba(OH)2
longer list FYI:
Lithium hydroxide 	LiOH
Sodium hydroxide 	NaOH
Potassium hydroxide 	KOH
Rubidium hydroxide 	RbOH
Cesium hydroxide 	CsOH
Magnesium hydroxide 	Mg(OH)
2
Calcium hydroxide 	Ca(OH)
2
Strontium hydroxide 	Sr(OH)
2
Barium hydroxide 	Ba(OH)
2
Tetramethylammonium hydroxide 	N(CH
3)
4OH
Guanidine 	HNC(NH
2)
2

note: bases tend to have a metal cation (and OH-) so many are ionic compounds rather than molecular (acids tend to have a nonmetal anion and H+)
ionic compounds dissociate in solution in proportion to how much they dissolve. if an ion dissolves completely in solution it’s a strong electrolyte

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104
Q

are all ionic compounds strong electrolytes?

A

no, because not all are soluble lol. the ones that are soluble are strong electrolytes – they dissociate 100% into ions.

for absolute accuracy: if it’s marked as a weak acid/base and it’s a salt it’s a weak electrolyte.
soluble ionic salts that are not marked as ‘weak’ are strong electrolytes.

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105
Q

when strong bases dissociate in water, what do they form?

A

cation and hydroxide ion

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106
Q

K2SO4 is a strong electrolyte.

Determine the concentration of each of the individual ions in a 0.400 M K2SO4 solution.

A

[K+]=.800M

[SO4[2-]]=.400M

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107
Q

molar mass – give measure

where do you find this value

A

g/mol

periodic table

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108
Q

M =

A

molarity

M = moles of solute / litre of solution

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109
Q

define hygroscopic

A

tending to absorb moisture from the air

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110
Q

Alisha wants to prepare a 0.310 M glucose solution using 1.50 moles of glucose. What volume of solution should Alisha prepare in liters, L?

SET UP THE EQUATION

A

1.50 mol glucose/0.310 M=4.84 L

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111
Q

define open, closed, and isolated systems

won’t hurt to give an example of each

A

open system: matter and energy can move between system and surroundings

  • fireplace with wood
  • open beaker

closed system: energy but not matter can move between system and surroundings

  • pressure cooker
  • flask with a cork

isolated system: neither matter nor energy can enter or leave system

  • thermos, approximately
  • flask with cork and high-quality insulation
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112
Q

what is a joule

A

1kg*m^2 / s^2

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113
Q

define w and q

A

w is work – energy resulting from a force acting on an object over a distance

q is heat – flow of energy that causes a temperature change

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114
Q

energy is

A

the capacity to do work or transfer heat

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115
Q

two forms of mechanical energy:

A

potential and kinetic

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116
Q

kinetic energy is

A

energy of motion

energy attributable to an object’s motion

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117
Q

potential energy is

A

energy of position

energy attributable to an object’s position; energy resulting from an object’s position or state

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118
Q

kinetic energy equation

A

KE = (1/2) mv^2

half of the mass times the squared velocity

119
Q

SI unit of energy?

A

joule

asked in another card:
1kg*m^2 / s^2

120
Q

A calorie is defined as

A

the amount of energy needed to raise the temperature of exactly 1 g of WATER by exactly 1°C.

121
Q

how many joules per calorie?

A

4.184 J in 1 cal

122
Q

first law of thermodynamics states

A

that the energy in the universe is constant

law of conservation of energy

123
Q

internal energy key point

give symbol used

A

U
internal energy, U, of a system is defined as the sum of all kinetic and potential energies of the PARTICLES within a system.

124
Q

energy can exchange between the system and surroundings in the forms of:

A

work and heat

ΔU = q + w

125
Q

positive q means

A

The system gains heat.

heat enters the system

126
Q

negative q means

A

The system loses heat.

heat leaves the system

127
Q

positive w means

A

Work is done ON the system.

work enters the system

128
Q

negative w means

A

Work is done BY the system.

work leaves the system

129
Q

ΔU positive means

ΔU negative means

A

The system gains internal energy.

The system loses internal energy.

130
Q

units for q and w?

A

J joule

131
Q

convert calories and Calories to joules

not necessary to memorise number

A

1 Cal = 1000 cal = 1 kcal = 4184 J
or
1 cal = 4.184 J

132
Q

is distance travelled a path or state function?

A

path function

(the state function would be net displacement or something)

If you are at home at 8:00 a.m. and also at home at 8:00 p.m., it does not necessarily mean that you have traveled zero miles that day.

133
Q

Classify each function as a path function or a state function:

work
enthalpy
heat 
energy 
distance travelled
A

path functions:
heat, work, distance travelled

state functions:
enthalpy, energy

134
Q

Work can be expressed as:

(Give expression if you remember it, but it’s on formula sheet)

𝑤=−𝑃Δ𝑉

or (for gases)

define the variables in the above equation(s)

A

Work can be expressed as
𝑤=−𝑃Δ𝑉

or
𝑤=−Δ𝑛𝑅𝑇
——–

𝑃 is pressure, 𝑉 is volume, 𝑛 is the number of moles, 𝑅 is the ideal gas constant, and 𝑇 is the temperature.

135
Q

Work can be expressed as
𝑤= −𝑃Δ𝑉
or
𝑤= −Δ𝑛𝑅𝑇

when dealing with gases, how would you know if a system had done work on its surroundings?

A

e.g., the number of moles of gas increases as the reaction proceeds.

The change in volume, Δ𝑉, and/or the change in the number of moles, Δ𝑛, are positive if there is an increase in the number of moles of gas as the reaction proceeds.

136
Q

read this again

At constant pressure, which of these systems do work on the surroundings?

A(g)+2B(g)⟶2C(g)
A(s)+B(s)⟶C(g)
2A(g)+2B(g)⟶5C(g)
A(s)+2B(g)⟶C(g)

A

in this case, w can be pictured as expansion, or the products having greater volume than the reactants

A system that does work on the surroundings has a negative value for work, 𝑤.

Work can be expressed as
𝑤=−𝑃Δ𝑉
or
𝑤=−Δ𝑛𝑅𝑇

where Δ𝑉
and Δ𝑛

are positive if there is an increase in the number of moles of gas as the reaction proceeds.

The other terms (𝑃,
𝑇, and 𝑅)

always have positive values.

These reactions

A(s)+B(s)⟶C(g)

2A(g)+2B(g)⟶5C(g)

produce more moles of gas than they consume, so Δ𝑉
and Δ𝑛 are both positive and 𝑤 is negative.

These reactions

A(g)+2B(g)⟶2C(g)

A(s)+2B(g)⟶C(g)

consume more moles of gas than they produce, and so Δ𝑉
and Δ𝑛 are both negative and 𝑤 is positive.

137
Q

internal energies of two substances are equal if (2 things)

A

if the MASSES and TEMPERATURES are identical.

138
Q

internal energy of a system is __________ the pathway taken to achieve it

A

independent of

139
Q

internal energy is a _______ function

A

state (as opposed to path)

140
Q

Path functions _________ the current state of a system. For example,

A

do not describe

systems are not said to possess a specific value of heat or work (q or w. ΔΗ is enthalpy)

141
Q

path or state function?

temp
pressure

A

state

state

142
Q

define pressure-volume work

A

The work done on or by a system when there is a volume change against an external pressure

143
Q

given the equation: 𝑤=−𝑃Δ𝑉

When the gas inside the flask increases in volume, the piston moves _____ and ΔV is a _____ value. The value of w in this case is _____, because the _____ is doing work on the ______.

If the volume of gas inside the flask DECREASES, on the other hand, then the piston moves ______ and ΔV has a _____ value. The value of w in this case is ______, because now the surroundings are doing work on the system.

A

When the gas inside the flask increases in volume, the piston moves UPWARDS and ΔV is a POSITIVE value. The value of w in this case is NEGATIVE, however, because the SYSTEM is doing work on the SURROUNDINGS.

If the volume of gas inside the flask decreases rather than increases, then the piston moves DOWNWARDS and ΔV has a NEGATIVE value. The value of w in this case is POSITIVE, because now the surroundings are doing work on the system.

144
Q

L · atm

is a unit of ______, equal to ______

A

energy

101.325 J. don’t need to memorise this value

145
Q

𝑤=−𝑃Δ𝑉

w will be reported in which units?

A

L · atm

146
Q

define enthalpy

A

The sum of internal energy of a system and the product of its pressure and volume change

147
Q

give enthalpy function (use H for enthalpy in this case)

identify what any symbols or letters represent

A

H = U + PV

enthalpy = internal energy + (pressure)(volume)

148
Q

At constant pressure, the CHANGE IN ENTHALPY (ΔH) is equal to

(may represent with an equation)

A

the change in internal energy plus the product of the pressure and change in volume.

ΔΗ = ΔU + PΔV

149
Q

ΔH is ______ when heat is transferred from the surroundings to the system.

(& is this endothermic or exothermic?)

A

positive

endothermic

150
Q

ΔH is ______ when heat is transferred from the system to the surroundings.
(endothermic or exothermic?)

A

negative

exothermic

151
Q

ideal gases are not subject to

A

interparticle interactions

152
Q

amount of heat required to change the temperature of a substance

give equation & be able to identify any variables

A

q = mcΔT

amount of heat required to change temp =
mass * specific heat of that substance (c) * change in temperature (C or K)

153
Q

q = mcΔT

this is the equation for:

change in temp is in which units

A

amt of heat required to change temp

either C or K will do because it’s only the net you’re looking for

154
Q

define specific heat

give proper units

identify specific heat in this equation:
q = mcΔT

A

The quantity of heat required to raise the temperature of 1 g of a substance by 1°C
(in absence of a phase change)

(J/g · °C)

c

155
Q

Substances with higher specific heat values resist temperature changes ____ than substances with lower specific heat values.

A

more

156
Q

put another way, substances with higher specific heat are more ______ to temp changes

A

resistant

157
Q

liquid water has an unusually ____ specific heat

A

high

(((it’s 4.184, twice as high as other-phase water which is nearly twice as high as some other common things; many metals are well below 1.0; only exception on the list is H2 with specific heat above 14)))

158
Q

specific heat units

A

J/ (g · °C)

159
Q

heat as a system reaches thermal equilibrium, write equation

A

(negative) - mass(A) times specific heat (A) times temp change (A) = (positive) mass(B) times specific heat (B) times temp change (B)
- mcΔΤ [thing giving heat] = mcΔT [thing gaining heat]

160
Q

if not given, assume the density of water to be

A

1g/mL

161
Q

calorimetry is

A

The study of heat transfer between substances by measuring the temperature changes of the substances involved

162
Q

To determine Δ𝐻rxn, as a STANDARD,

A

express 𝑞rxn (q is specific heat of solution x change in temp x total mass) in units of kilojoules per mol of X. That means you’ll need to convert the mass of X to moles using the molar mass, then divide the number of kilojoules by the number of moles to get kJ/mol.

163
Q

bomb calorimeter heat flow of reaction

q[rxn]=

A
  • C(cal)ΔT
164
Q

J joule can be the units for:

A

ΔU, q or w

165
Q

gas directly to a solid is called

A

deposition

166
Q

Wood is burned in an open fireplace. The reaction releases 250 kJ of heat and does 15.0 kJ of work on the surroundings.

Determine whether this reaction is endothermic or exothermic.

Calculate ΔH for this reaction.

Calculate ΔU for this reaction.

if you need a hint this is on your equation sheet:
ΔH = ΔU + PΔV

A

Heat is transferred from the system to the surroundings, so the reaction is exothermic (i.e., ΔH < 0).

Since this is an open system, the pressure is constant. Therefore, ΔH is equal to the amount of heat transferred and it has a negative value because the heat is transferred from the system to the surroundings: ΔH = q = −250 kJ.

ΔH depends only on q, but ΔU depends on q and w. In this case, the work (15.0 kJ) is done on the surroundings, so it also has a negative value.
-265 kJ

167
Q

what part of the kinetic energy equation do you continually forget

A

KE = (1/2) mv^2

ONE HALF

168
Q

watch your + / - signs

Calculate the final temperature after 6.80 g of aluminum (c = 0.892 J/g · °C) at 67.4°C is placed in 195 g of water at 19.6°C.
–set up the equation, just using symbols no #s

A

q = 0 = {metal}mcΔT+{water}mcΔT

0 = mcΔT{metal} + mcΔT{water}

169
Q

When heat is transferred from one object to another in a controlled experiment, and no heat is transferred to or from the surroundings, the heat gained or lost by the system is equal to

A

zero.

170
Q

Constant-pressure calorimetry uses atmospheric pressure and relies on the following equation.

q[sub]p =

(For the sake of simplicity, the constant-pressure calorimetry example does not account for absorption of heat by the calorimeter walls.)

In constant-pressure calorimetry, the heat absorbed or released by the solution, q[soln], has the _____ sign of the HEAT ABSORBED OR RELEASED BY THE REACTION:

????

A

q[sub]p = ΔH

opposite
q[p]

(((in other words:
q[sub]p = – q[soln] or vice versa

171
Q

The release or absorption of heat that occurs when a solid compound, such as KCl or CaCl2, dissolves in a solvent at constant pressure to infinite dilution is known as the _________________

A

enthalpy of solution, ΔH[soln].

172
Q

water’s density varies with temperature.

however, if you have to guess it, try this:

A

1.0 g mL^−1
aka
1g/mL

173
Q

Calculate the enthalpy change of a reaction (pressure calorimetry)

what do you need to remember in reaction enthalpies?

in an exothermic reaction, ΔH is _______
in an endothermic reaction, ΔH is _______

A

ΔH = q[p] = –q[soln]

use the correct +/- sign!

negative
positive
ALWAYS!

174
Q

calculate enthalpy change per mole of acid (2 steps, each requires you to know or have calculated a certain value in advance)

what unit will the answer have?

A

Calculate the number of moles of acid
–>molarity required
Divide the enthalpy change of the reaction by the number of moles of acid.
–>calculate enthalpy change of the reaction.

J/mol, preferably kJ per mol

175
Q

convert cm^3 to mL

A

1 cm^3 = 1 mL

176
Q

constant-volume calorimeter vs constant-pressure calorimeter

A

constant-volume calorimeter
bomb calorimeter
more accurate d/t better insulation (truly an isolated system)
heat transferred corresponds to the change in internal energy, ΔU. Recall that ΔU = q + w and that the work term, w, is zero inside a rigid container with a constant volume.
use the total heat capacity of the calorimeter components and the water inside the calorimeter, rather than just the specific heat of the water (aka C[cal])
–>Once C[cal] is known, it can be used in all reactions for that particular calorimeter with that particular volume of water.

constant-pressure calorimeter
useful for students. thermos inside a thermos
q[p] = ΔH
q[soln] = –q[p]
–q[soln] = q[p] = ΔH
use specific heat of water for calculations

177
Q

inside a rigid container with a constant volume, the work term (w) is:

A

zero

178
Q

bomb calorimeter

is it constant-volume or constant-pressure

A

constant-volume

179
Q

C[cal] indicates ______

in which system?

A

total heat capacity / combined heat capacity
(of calorimetre components and water)
constant-volume or bomb calorimetry

180
Q

q[p] and q[v] indicate

A

q[p] is the heat flow inside a constant-pressure calorimeter

q[v] is heat flow inside a constant-volume calorimeter

181
Q

q[soln] is different in constant-volume vs constant-pressure calorimetry. describe

A

constant pressure: think of ‘system’ as water

constant volume (bomb): think of ‘system’ as the machine. it’s all the components adjusted to a particular volume of water. water is included

182
Q

in bomb calorimetry, q[soln] =

it is _____ in sign to the heat flow of the reaction

A

q[soln] = C[cal]ΔT

opposite
q[soln] = –q[v]

183
Q

give the basic equation(s) for bomb calorimetry

A 
q[soln] =       C[cal]ΔT
q[v]      =     --C[cal]ΔT

ΔU[rxn] = q[v] = –C[cal]ΔT

184
Q

if ΔH < 0, the reaction is ___________
if ΔH > 0, the reaction is ___________

– ΔH has the same sign as q[___]

q[p] or q[v] = ΔU[___]

A

exothermic
endothermic
ALWAYS

ΔH describes reaction, q[p] or q[v]
(((ΔH = q[p] …. ΔH does not equal q[v] necessarily because while the volume is constant the pressure probably isn’t)))

q[p] or q[v] = the ΔU reaction

185
Q

mcΔT, what do you have to remember about mass

A

sum up the mass of reactants

186
Q

energy released, ΔH ______

A

negative

187
Q

Given the thermochemical equations

A(g)⟶B(g)Δ𝐻=80 kJ
B(g)⟶C(g)Δ𝐻=−150 kJ

find the enthalpy changes for each reaction.

3A(g)⟶3B(g)Δ𝐻=
B(g)⟶A(g)Δ𝐻=
A(g)⟶C(g)Δ𝐻=

A

Δ𝐻=

240 kJ

  • 80 kJ
  • 70 kJ
188
Q

constant volume calorimeter

identify system and surroundings

A

system: reaction
surroundings: solution and calorimeter including all components

189
Q

standard conditions

A

25C, 1 atm pressure

190
Q

standard enthalpy of formation notation

Δ𝐻[subscript: f][superscript °]
what do these little notations mean

A

Δ𝐻 is enthalpy
[subscript: f] – formation reaction
[superscript °] – standard state (not stp) (25C, 1atm)

191
Q

describe standard enthalpy of formation

A

enthalpy of formation of 1 mol compound from its constituent elements in their standard states

192
Q

standard enthalpy of formation – Δ𝐻[f]° – of any ELEMENT in its standard state is

A

0

doesn’t take any energy to form an element. presumably it already exists.

H2(g), Fe(s), Cl2(g), Na(s), O2(g)
if it’s diatomic or in a certain phase at 25C 1atm, that’s the standard state

193
Q

explain how to apply hess’ law

A
  1. set up equations so products and reactants are on the correct sides. if you switch the direction of the reaction, switch the +/- for that reaction’s Δ𝐻.
  2. cross out anything that appears on either side. get to the target reaction. multiply across if you need, but make sure to multiply that reaction’s Δ𝐻 by the same coefficient
  3. ADD Δ𝐻 values together for net or total
194
Q

standard enthalpy of reaction (Δ𝐻°[rxn]) =

A

Δ𝐻°[f] products - Δ𝐻°[f] reactants

195
Q

define bond enthalpy

A

amt of energy required to break a specific bond of 1 mol (gaseous molecules)

196
Q

the ______________ within a molecule determines the stability of the molecule

A

strength of the bonds

197
Q

are high bond enthalpies stable or unstable?

A

high bond enthalpies – bonds very stable. you have to put a lot of energy in to break them to make a reaction occur

198
Q

using bond enthalpies, predict which is more reactive, HCl with a bond enthalpy of 432kj/mol, or HBr, with a bond enthalpy of 366kj/mol?

Explain

A

HBr.

takes more energy to break the bond in HCl, since bond enthalpy is higher

199
Q

forming bonds _______ energy

breaking bonds _________ energy

A

forming bonds gives off energy

breaking bonds uses energy

200
Q

forming bonds _______ energy

breaking bonds _________ energy

A

forming bonds releases energy

breaking bonds absorbs energy

201
Q

forming bonds releases energy
breaking bonds absorbs energy

now describe endothermic and exothermic reactions

A

endothermic: breaking bonds absorbs MORE energy than forming bonds releases
exothermic: forming bonds releases MORE energy than is absorbed breaking bonds

202
Q

as bond length increases, bond enthalpy _____

A

decreases

longer bonds are weaker bonds, more shielding going on
recall the size trend on the periodic table, towards grp1 and further down the molecules are larger, exerting a weaker pull on their valence electrons

shorter bond, harder to separate those atoms

203
Q

breaking bonds altogether, breaking all bonds,

which is the easiest to break?

C - = C (triple bond lol)
C = C
C - C

A

breakig the pi bond on a carbon carbon triple bond is the easiest but as far as bond enthalpy, breaking all the bonds, it’ll be easiest to break just the c-c single bond

204
Q

give bond orders for:
single bonds
double bonds
triple bonds

A

bo = 1
2
3

205
Q 
[lab]
titration
SA + SB --> salt + water
flask acid, burette base
colour change =
A

mol H+ = mol OH-

experimental approximation of equivalence point

206
Q

to get from gas to liquid, ______ heat

A

remove

207
Q

pressure is

A

force per unit area

consider all gas particles and divide by area of surface

208
Q

enthalpy is

A

heat flow at constant pressure

209
Q

Δ𝐻[soln] is what?

if the solution gets hot, what is the sign for Δ𝐻

A

enthalpy of a dissolution reaction

if the solution gets hot, the reaction was exothermic, and so the reaction gave off energy. Δ𝐻 should therefore be negative.

210
Q

dissolving NaCl in H2O makes the water colder. explain

will this spontaneously reverse ?

A

endothermic reaction

requires heat for the NaCl to dissolve in water

no it doesn’t spontaneously fall out of solution

211
Q

explain the concept of thermal equilibrium

use an example of a piece of hot metal placed in a cup of water

A

when the system reaches thermal equilibrium, the metal is giving heat to the water at the same rate as the water is giving heat to the metal

q[1] = -q[2] – doesn’t matter on which side you put the negative

212
Q

given density of solution at 1.00 g/mL

50 mL of solution should have which mass

A

50g

213
Q

Hess’ law

When two or more processes combine to give a resulting process, their enthalpy changes add to give the enthalpy change for the resulting process

There are three rules but give 4

A

(1) When an equation is reversed, the sign of its enthalpy changes.
(2) when the coefficients in an equation are multiplied or divided by a factor, the enthalpy value is multiplied or divided by that same factor.
(3) When reactions are summed, the enthalpy of the overall reaction is the sum of the enthalpies of the component reactions.
4. cross out identical items appearing on either side of the –> to simplify

214
Q

standard state of a substance

2 considerations

there’s one additional consideration for carbon e.g. (is it graphite or diamond lol)

A

diatomic or not

phase (solid liquid gas)

215
Q

which elements are diatomic at STP

A

have no fear of ice cold beer

also, H —–> NO (top 4 halogens)

H
N
F
O
I
Cl
Br
216
Q

STP is NOT the standard state we’re using! use this _________

Until 1982, STP was defined as a temperature of ______ and an absolute pressure of ______

[[after 1982, STP was changed to ______]]

A

298(.15) K (25.00 °C) and 1 atm

273.15 K (0 °C, 32 °F)

exactly 1 atm (101.325 kPa).

[[273.15 K and 100 kPa (1 bar)]]

217
Q

standard enthalpy of formation is:

A

the enthalpy of formation
of 1 mol of a compound
from its constituent elements in their standard states
which for purposes of this class are 25C and 1atm

218
Q

standard enthalpy change for any reaction =

A

Δ𝐻°[rxn] = ΣΔ𝐻°[f]{products} – ΣΔ𝐻°[f]{reactants}

standard enthalpy change for any reaction = sum of the standard enthalpies of formation of all products minus the sum of the standard enthalpies of formation of all reactants

219
Q

standard enthalpy change for any reaction in cases where there are multiple moles of reactants and products =

A

Δ𝐻°[rxn] = ΣnΔ𝐻°[f]{products} – ΣnΔ𝐻°[f]{reactants}

n= number of mols of product / reactant from balanced chem equation

220
Q

once you have the standard enthalpy change for a reaction (Δ𝐻°[rxn]), can you easily find the standard enthalpy change for, for instance, 2.5 mol of the given substance ?

A

multiply Δ𝐻°[rxn=1 mol] by 2.5 or whatever given mol number

221
Q

if you get a result Δ𝐻°[rxn] = -5kJ, and it asks you how much heat was released, how should you report it?

A

5kJ heat is released (Δ𝐻°[rxn]= –5kJ)

222
Q

what is the molar mass of hydrogen at XTP (alternative STP with T=25C) – you idiot.

A

2.016g/mol, you idiot. it’s diatomic

223
Q

Bond strengths are generally measured as bond enthalpies, which are defined as

A

the enthalpy change, ∆H (Section 10.4), associated with breaking a specific bond in 1 mol of gaseous molecules.

224
Q

Energy must be ______ to break bonds; bond breaking is ______.

Energy must be ______ when bonds are formed; bond making is ______.

A

added
endothermic

released
exothermic

225
Q

Which is more reactive, AQ4, with avg bond enthalpies for AQ (with appropriate amt of σ and π bonds) at 410kJ per mol, or AY4, with avg bond enthalpies for AY (appropriate bonds) at 330 kJ/mol?

A

AY4 is more reactive because it takes less energy to break the AY bonds

226
Q

Bond enthalpy values are often used to estimate enthalpies for reactions involving _______

A

molecular compounds

227
Q

if bond enthalpy is higher, it takes _____ energy to break the bonds, and the compound formed using that bond enthalpy (with all else being the same) is _____ reactive

A

more

less

228
Q

Using bond enthalpies to estimate reaction enthalpy is an example of Hess’s law. The process is:
(4 steps)

A

ENSURE REACTION IS BALANCED
Calculate the total energy that would be added to turn the reactant molecules into individual atoms (bond-breaking is endothermic),
Calculate the total energy that would be released when these atoms combine to form the product molecules (bond-making is exothermic), and
Compare the two energy values:

Δ𝐻[rxn] = Σ(Δ𝐻 of bonds broken) - Σ(Δ𝐻 of bonds formed)

229
Q

bond-breaking is ______

bond-making is ______

A

endothermic

exothermic

230
Q

Bond enthalpies –> est REACTION ENTHALPY
reactant molecules –> atoms: calculate energy _____
atoms –> products: calculate energy _____
compare in this equation:

A

*make sure to balance original equation first

added
released

Δ𝐻[rxn] = Σ(Δ𝐻 of bonds broken) - Σ(Δ𝐻 of bonds formed)

231
Q

If it takes more energy to break the bonds in the reactant molecules than is released when making the bonds in the product molecules, then the reaction must ______ and, therefore, is _____. If more energy is ______ in bond-making than was used in bond-breaking, on the other hand, then the reaction is ______

A

absorb energy
endothermic

released
exothermic

232
Q

estimating ∆Hrxn, final answer, –690 kJ, indicates

A

a highly exothermic reaction, which is why methane (natural gas) is used as a fuel.

233
Q

Published bond lengths are ____ values obtained from the measurements of many different molecules that contain bonds between the same pairs of atoms, just as bond enthalpies are.

A

average

234
Q

Describe the relationship between bond lengths and bond enthalpies for single, double, and triple bonds between carbon atoms.

A

In going from single bond to double bond to triple bond, bond length decreases while bond enthalpy increases.

235
Q

For similar atom pairs, there is _____ relationship between bond length and bond enthalpy

A

an inverse

236
Q

shorter bonds are ____ bonds.

A

stronger

237
Q

compare standard enthalpies of formation (equation) to that of bond enthalpies. what is strange? is one different from the patterns we’ve generally seen?

A

formation: products minus reactants
Δ𝐻rxn=Δ𝐻𝑓[products]−Δ𝐻𝑓[reactants]

bond enthalpies: reactants minus products
Δ𝐻rxn=Δ𝐻bond[reactants]−Δ𝐻bond[products]

238
Q

Given that

Δ𝐻∘f[Br(g)]=111.9 kJ⋅mol−1

Δ𝐻∘f[C(g)]=716.7 kJ⋅mol−1

Δ𝐻∘f[CBr4(g)]=29.4 kJ⋅mol−1

calculate the average molar bond enthalpy of the carbon–bromine bond in a CBr4

molecule.
Δ𝐻C−Br=

A

Write the balanced chemical equation for the breaking of all bonds in CBr4

without the formation of any new bonds.

–For example, the breaking of all bonds in H2Ois written as
H2O⟶2H+O

The enthalpy of that reaction, Δ𝐻rxn,
can be expressed in two ways: (1) in terms of enthalpy of formation, Δ𝐻f, and (2) in terms of bond enthalpy, Δ𝐻bond.

Δ𝐻rxn=Δ𝐻𝑓[products]−Δ𝐻𝑓[reactants]

Δ𝐻rxn=Δ𝐻bond[reactants]−Δ𝐻bond[products]

Calculate Δ𝐻rxn
using the first formula, then plug that value into the second formula. If you properly wrote the chemical equation, Δ𝐻bond[products] will be equal to zero. Be sure to take into account the stoichiometric coefficients.

239
Q

Write Δ𝐻rxn in terms of enthalpy of formation, Δ𝐻f,

A

Δ𝐻rxn=Δ𝐻𝑓[products]−Δ𝐻𝑓[reactants]

240
Q

You have not correctly predicted the energies of the bonds in NO2. The Lewis structure of NO2 shows one single bond and one double bond, but the true structure is a resonance hybrid.

What are the bond energies?

A

an average of the single and double bond, really

241
Q

The stability of ionic compounds is mainly attributed to the

A

electrostatic attraction between ions of opposite charge

ionisation energy and electron affinity

242
Q

lattice energy define

give the symbolic indicator

A

energy required to separate an ionic compound into its constituent gaseous ions

a couple more definitions loL:
Energy is released when ions assemble into a lattice structure
(energy released when gas-phase ions are converted into a solid ionic compound)

Δ𝐻[L]

243
Q

define bond energy

A

The energy needed to break a mole of a specific type of bond; also called bond enthalpy

244
Q

what’s the Born-Haber cycle?

A

A thermodynamic cycle that uses Hess’s law, ionization energy, electron affinity, and the energies of other processes to calculate the lattice energy of an ionic compound

245
Q

Lattice energy: give steps
KCl is the example

write equation

A
  1. reverse enthalpy of formation of compound
    *—Δ𝐻[f] of KCl:
    KCl(s) ⟶ K(s) + ½ Cl2(g)
  2. conversion of K(s) to K(g) (enthalpy of sublimation, Δ𝐻[sub]):
    K(s) ⟶ K(g)
    **because XTP of potassium is K(s); all must be in gaseous form
  3. ionization of K(g) to K+(g):
    K(g) ⟶ K+(g) + e–
    ((Δ𝐻 = IE))
    **because all must be ions
  4. formation of Cl(g) from Cl2(g) (bond energy (Δ𝐻 = BE) of Cl2):
    Δ𝐻 [Cl-Cl]
    **because ratio of K to Cl is 1:1, isolate Cl
  5. addition of an electron to Cl(g) to form Cl–(g), the electron affinity, EA, of chlorine:
    Cl(g) + e– ⟶ Cl–(g)
    ((Δ𝐻 = EA))

Δ𝐻[L] = — ( — Δ𝐻[f] + Δ𝐻[sub] + IE + BE + EA)

246
Q

ionic radius ______ moving down a group in the periodic table

as you move down a group therefore, lattice energies _____ in magnitude as sizes of cations and anions increase

as lattice energies increase, potential energies _____

In other words, the formation of ionic compounds becomes ____ exothermic as the size of the ions increase.

This is because

A

increases

decrease

increase

less

larger ions means longer ionic bond lengths (electrostatic attraction gets weaker as nuclei move further apart)

247
Q

write equation representing the potential energy of two interacting charged particles, E[el] as it relates to the distance between the centre of the two particles, d.

(On formula sheet)

A

E[el] = k (q1q2) / d

q1 and q2 are the charges of the particles and k is the constant 8.99 × 109 J·m/C2.

the potential energy of two interacting charged particles, E[el], is inversely proportional to the distance between the center of the two particles, d.

248
Q

The coulomb, C, is the derived SI unit for

A

electrical charge.

249
Q

Lattice energies tend to be _____ in magnitude (more _____) for ionic compounds consisting of ions with large charges.

A

larger in magnitude / more negative

250
Q

The bond length of MX (1 metal, 1 halogen) is determined by:

A

summing the ionic radii of its ions

251
Q

lattice energy depends on: (give 2 things)

A 
bond length (longer bonds, less energy)
charges of the ions (stronger charges, more energy)
252
Q

MgO and LiF have nearly the same bond lengths. which of these two ionic compounds will have a more negative lattice energy?

A

MgO

E[el] = k (q1q2) / d
q1 and q2 are the charges of the particles and k is the constant 8.99 × 109 J·m/C2.

k is a constant, d is nearly the same (given in the problem)
LiF q1 * q2 = - 1 * 1 = - 1
MgO q1 * q2 = - 2 * 2 = - 4

253
Q

define monovalent & divalent ions

name a trivalent ion

A

Monovalent cations lose 1 valence electron (+1)

Divalent cations lose two valence electrons and attain a 2+ charge.

Aluminium (3+)

254
Q

Endothermic reaction has _____ reactant bonds and _____ product bonds

A

Endothermic reaction has stronger reactant bonds (large energy requirement) and weaker product bonds (small energy release).

255
Q

exothermic or endothermic?

A–B bonds are stronger than A–C bonds.

The reaction AB+C⟶AC+B is:

A

endothermic

256
Q

exothermic or endothermic?

A–A and C–C bonds are both stronger than A–C bonds.

The reaction A2+C2⟶2AC is

A

endothermic

257
Q

exothermic or endothermic?

B–B and C–C bonds are both stronger than B–C bonds.

The reaction B2+C2⟶2BC is

A

endothermic

258
Q

exothermic or endothermic?

The energy required to break the reactant bonds is less than the energy released during the formation of product bonds.

A

exothermic

In other words, an exothermic reaction has weaker reactant bonds (small energy requirement) and stronger product bonds (large energy release).

259
Q

exothermic or endothermic?

B–C bonds are weaker than A–B bonds

The reaction A+BC⟶AB+C is

A

exothermic

260
Q

exothermic or endothermic?

A–A and B–B bonds are both weaker than A–B bonds.

The reaction A2+B2⟶2AB is

A

exothermic

261
Q

density of gas phase ____ density of liquid phase ____ density of solid phase

A

is less than
is approximately equal to

Despite some variation between the densities of the solid and liquid phases of a given substance, it is usually negligible compared to the drastic drop in density seen in the transition to the gas phase.

262
Q

conversions you must memorise:

atm vs mmHg vs torr

A

1 atm = 760 mmHg

1 mmHg = 1 torr

1 atm = 760 torr

263
Q

which measures of temp can you use in gas laws?

A

ONLY KELVIN

264
Q

STP (or as i’ve begun to call it due to variability, XTP) in thermochem vs gas laws?

A

thermochem:
25C is xtp
K or C are fine for relative calculations. C is common

gas laws:
0c is xtp
use 1 atm but it’s officially .987atm
gas chem, temp is in K!!!

265
Q

which two pressure measurements are equivalent?

A

mmHg and torr

the two with doubled letters

266
Q

Write equation for Boyle’s law. (More info below.)

What units are used?

The scientist Robert Boyle found that the volume of a gas varies inversely with the pressure exerted by the gas, provided the number of moles and temperature of the gas are held constant. This relationship is known as Boyle’s law. Stated a different way, Boyle’s law says that the product of pressure (𝑃) and volume (𝑉) is a constant.

Boyle’s law is often used to calculate the volume resulting from a pressure change and vice versa. The product of the initial pressure (𝑃1)
and the initial volume (𝑉1) is equal to the product of the final pressure (𝑃2) and the final volume (𝑉2).

A

𝑃1𝑉1=𝑃2𝑉2

L and atm are fine.

as long as you’re doing proportions, just keep it consistent. e.g., mL to mL, fine

267
Q

Write equation for Charles’ law. (More info below.)

What units are used?

Charles’s law can be used to relate the initial volume and temperature of a gas to the final volume and temperature of a gas when the number of moles and pressure remain constant.

The scientist Jacques Charles found that the volume of a gas varies directly with the absolute temperature (temperature in units of kelvins), provided the pressure and number of moles of gas are held constant. This relationship is known as Charles’s law. Put another way, Charles’s law states that the ratio of volume (𝑉) and temperature (𝑇) is a constant.

Charles’s law is often used to calculate the volume resulting from a temperature change and vice versa. The ratio of the initial volume (𝑉1)
and initial temperature (𝑇1) is equal to the ratio of the final volume (𝑉2) and final temperature (𝑇2).

A

𝑉1 / 𝑇1 = 𝑉2 / 𝑇2

Kelvin
L, mL. proportions, just be consistent

268
Q

more pressure conversions, yay. this one’s more complete. atm, torr, Pa or kPa, mmHg

A

The millimeter of mercury, mmHg, was one of the first units used for measuring pressure. It came into use because many instruments measured pressure based on the height of the mercury column the gas can support. The unit mmHg is also known as the torr, named after Torricelli, the inventor of the first barometer. A related unit is the standard atmosphere, atm, which is approximately equal to the atmospheric pressure at sea level. The relationship between these three pressure units is

760 mmHg=760 Torr=1 atm

The SI unit of pressure is the pascal, Pa, where

1 atm=101325 Pa

269
Q

combined gas law – give equation

A

𝑃1𝑉1 / 𝑇1 = 𝑃2𝑉2 / 𝑇2

formatting!

p1v1 = p2v2
___ ____
t1 t2

270
Q

Write equation for Avogadro’s law. (More info below.)

What units are used?

The relationship between the volume and the number of moles of gas is known as Avogadro’s law. Avogadro’s law states that the volume of a gas varies directly with the number of moles of gas, provided the pressure and temperature of the gas are constant. Put another way, Avogadro’s law states that the ratio of volume (𝑉) and number of moles of gas (𝑛) is a constant.

Avogadro’s law is often used to calculate the volume resulting from a change in the number of moles of gas and vice versa. The most useful version of Avogadro’s law is the one that equates the ratios of the initial and final volume and number of moles,

A

𝑉1 / 𝑛1 = 𝑉2 / 𝑛2

mol and L are standard
ratios again, just be consistent

271
Q

Direct proportionality means that two quantities scale in the same way. Meaning, an increase in one quantity causes the other to increase. Similarly, a decrease in one quantity causes the other to decrease.

A

Inverse proportionality means that two quantities scale in the opposite way. Meaning, an increase in one quantity causes the other to decrease and a decrease in one quantity causes the other to increase.

272
Q

IDEAL gas law – give equation

what are the units ?

what is the constant ?

A

𝑃𝑉=𝑛𝑅𝑇

n - mol
V - L
P - atm
T - K!!!

𝑅 is a constant equal to 0.08206 L⋅atm/(mol⋅K)

273
Q

given density, temp, and pressure, determine molar mass of a sample

A

rearrange equation PV = nRT

n = PV / RT
n = mol; molar mass = g/mol
molar mass = m/n = mRT/pV
density = m/V
molar mass = densityRT/pressure
274
Q

according to the kinetic molecular theory of gases

the average kinetic energy of the gas particles is _________ to the ______ temperature of the gas.

A

directly proportional / absolute

The average kinetic energy of the gas particles is directly proportional to the absolute temperature of the gas

275
Q

The five postulates of the kinetic molecular theory explain why ideal gases behave as they do: (give all 5)

A
  1. gases = particles in RANDOM MOTION
  2. gas molecules have INSIGNIFICANT VOLUME compared with volume occupied
  3. COLLISIONS comprise the significant portion of intermolecular interaction. the particles move in constant straight-line motion unless they collide with other particles or the walls of the container,
  4. collisions are PERFECTLY ELASTIC: no energy is lost when molecules collide
  5. AVG KINETIC ENERGY directly proportional to ABSOLUTE TEMP

(now these, per text):

  1. Gases are composed of small molecules that are in constant, random motion.
  2. The volume that is taken up by the molecules themselves is insignificant compared with the overall volume occupied by the gas.
  3. Forces between the molecules are negligible, except when the molecules collide with one another.
  4. Molecular collisions are perfectly elastic; that is, no energy is lost when the molecules collide. (No friction e.g.)
  5. The average kinetic energy of the gas molecules is directly proportional to the absolute temperature of the gas. (the kelvin temperature of a gas (or of a liquid or solid as well) is a measure of the average kinetic energy of the particles that make up the sample.)
276
Q

what causes gases to exhibit pressure?

A

the particles of gases are in constant motion and are colliding with each other. The particles must also be colliding with the walls of their containers. The force of those collisions is source of gas pressure

277
Q

When the volume of a gas sample decreases without a change in the number of particles, the rate of collisions of those particles with the container walls must _____, causing ______ to increase.

A

increase

pressure

278
Q

an increase in temperature causes gas particles to have more energy and move at higher speeds. When the temperature of a gas is increased, these faster-moving particles collide with the walls of the container more frequently and with greater force, thus increasing the pressure. However, if the container is flexible:

A

the pressure will remain constant and the volume will increase to accommodate the faster-moving particles as predicted by Charles’s law

279
Q

When the number of moles of a gas are increased, the rate of collisions of those particles with the walls of the container ______

A

increases (explaining Avogadro’s law)

280
Q

Kinetic molecular theory tells you to consider only the number of particles of the gas and not the identity of the gas, which explains why…

A

the partial pressure of a gas in a mixture is related to the number of moles of each component gas (Dalton’s law)

281
Q

(2 things) are associated with deviations from the ideal gas law.

A

Very high pressures and very low temperatures

282
Q

The temperature of a gas is a measure of

The pressure of a gas results from

A

the kinetic energy of the gas particles.

gas particles colliding with the container walls.

283
Q

The average kinetic energy of the molecules in a gas sample depends only on the temperature, 𝑇.

However, given the same kinetic energies, a lighter molecule will move _____ than a heavier molecule, as shown in the equation for rms speed

rms speed= √ (3𝑅𝑇
‾‾‾‾‾
ℳ)

where 𝑅=8.314 J/(mol⋅K)
and ℳ is molar mass in KILOGRAMS per mole.

A joule is the same as a ___________

if you can fill in the two blanks, mark this question as acceptable

A

faster

a lighter molecule will move faster than heavier molecule given same kinetic energy

The root-mean-square speed is the measure of the speed of particles in a gas

rms speed = square root of the average velocity-squared of the molecules in a gas

The root-mean-square speed takes into account both molecular weight and temperature, two factors that directly affect the kinetic energy of a material and is represented as V rms = (sqrt ((3* [R] * T)/ M)) or Root mean square velocity = (sqrt ((3* [R] * Temperature of Gas)/ Molar Mass)).

A joule is the same as a kilogram‑meter squared per second squared (kg·m2/s2).

284
Q

molecular weight?

A

molar mass

285
Q

The average kinetic energy of the molecules in a gas sample depends only on the ______

However, given the same kinetic energies, a lighter molecule will move _____ than a heavier molecule, as shown in the equation for rms speed

A

temperature, 𝑇.

faster

286
Q

define diffusion

define effusion

A

diffusion:
The spreading out of gas particles by random motion and collisions to occupy an entire volume
(a gas sample is introduced into a larger volume and the gas particles spread out to occupy the entire volume, mixing with the other gases present)

effusion:
The movement of gas particles through a tiny opening without collisions
the escape of gas molecules through a tiny hole into a vacuum, is simpler to quantify than diffusion because it does not depend on collisions.

287
Q

According to Graham’s law, effusion rate is

A

inversely proportional to the square root of molar mass.

rate A / rate B = sqrt(molar mass B / molar mass A)

288
Q

average atomic mass of an element with two isotopes?

A

avg atomic mass = (fractional abundance isotope 1 x atomic mass isotope 1) + (fractional abundance isotope 2 x atomic mass isotope 2)

289
Q

define:

Ionisation energy

Electron affinity

A

ionisation energy:
energy required to remove an electron

electron affinity:
energy change when gaining an electron

290
Q

what does a more negative EA imply?

A

more negative: more favourable (more exothermic)

291
Q

electron affinity vs electronegativity

describe where they can be found

A

electron affinity – basically the group on the periodic table

electronegativity – special chart. F>O>Cl etc

292
Q

standard enthalpy of formation

given standard enthalpies of formation of reactants and products

A

products - reactants (multiply through mols and sum)

293
Q

effective nuclear charge (Z[eff])

A

Z-S
Z is atomic number
S is shielding
S = .85 (core electrons) + .35 (other val e- if any)

Z = #protons - .85(#core e-) - .35(#other vale-)

294
Q

Cations are _____ than their corresponding neutral atoms, whereas anions are _____ than their corresponding neutral atoms.

A

smaller

larger

汉语一级 (12 decks)

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