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AECHE > cap 5 - semester 2 topic test > Flashcards

cap 5 - semester 2 topic test Flashcards

acids and bases, stoichiometry, organic chemistry, reactions of hydrocarbons, energy and enthalpy, rates of reaction

1
Q

acids

A

substances that ionise to release H+ ions when dissolved in water

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2
Q

bases

A

substances that produce OH- ions in solutions

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3
Q

properties of acids:
pH, solutions, taste, feel

A

pH < 7, forms H+ ions, sour, stinging/burning skins

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4
Q

properties of bases:
pH, solutions, taste, feel

A

pH > 7, forms OH- ions, bitter, slippery feel

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5
Q

indicators

A

chemicals that change colour in the presence of an acid and alkali

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6
Q

indicating acids

A

blue litmus paper becomes red, OR
red/orange with universal indicator

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7
Q

indicating bases

A

red litmus paper becomes blue, OR
blue/purple with universal indicator

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8
Q

ionisation

A

where a covalent molecule splits into ions (i.e. all acids)

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9
Q

dissociation

A

where an ionic compound dissolves in water

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10
Q

strong vs weak

A

strong: completely ionise/dissociate in water
weak: only partially ionise/dissociate

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11
Q

HCl

A

strong acid

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12
Q

HNO3

A

strong acid

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13
Q

H2SO4

A

strong acid

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14
Q

CH3COOH

A

weak acid

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15
Q

H3PO4

A

weak acid

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16
Q

H2CO3

A

weak acid

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17
Q

common strong bases

A

group 1 and 2 oxides and hydroxides

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18
Q

NH3

A

weak base

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19
Q

Na2CO3

A

weak base

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20
Q

NaHCO3

A

weak base

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21
Q

concentrated

A

high molarity

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22
Q

diluted

A

low molarity

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23
Q

diprotic acids

A

acids that have two H+ in the structure and release both in solution e.g. H2SO4
(strongly lose first, weakly lose second)

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24
Q

pH scale

A

inverse logarithmic measure of hydrogen ion concentration

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25
Q

acid + metal

A

salt + hydrogen gas

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26
Q

acid + metal oxide

A

salt + water

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27
Q

acid + metal hydroxide

A

salt + water

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28
Q

acid + metal carbonate

A

salt + water + carbon dioxide

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29
Q

acid + metal hydrogencarbonate

A

salt + water + carbon dioxide

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30
Q

base + ammonium salt

A

salt + water + ammonia (pungent gas)

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31
Q

base + nonmetal oxide

A

salt + water

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32
Q

self-ionisation (autoionisation) of water

A

equilibrium reactions that water undergoes on its own, by ionising and recombining constantly

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33
Q

ionic product of water (Kw)

A

equillibrium constant for the self-ionisation of water

Kw = [H+]*[OH-] = 1 x 10^-14 at 25deg

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34
Q

neutrality when

A

concentration of hydrogen = hydroxide

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35
Q

why is carbon able to form such a diverse range of compounds?

A

2, 4 electron configuration (4 valence) ⇒ ability to covalently form 4 more single bonds ⇒ variety of arrangements (single, double, triple) ⇒ forms many compounds

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36
Q

hydrocarbons

A

organic compounds, based on hydrogen and carbon atoms

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37
Q

aliphatic

A

carbon atoms arranged in open chains ⇒ can be saturated or unsaturated

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38
Q

aromatic

A

carbon atoms arranged in closed rings

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39
Q

functional groups

A

atom or group of atoms responsible for typical chemical reactions, determining patterns of reactivity

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40
Q

homologous series

A

group of molecules with same functional group, but different number of carbons in the main chain

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41
Q

alkanes

A

single bonds only
saturated compounds
C(n)H(2n+2)

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42
Q

alkenes

A

at least one double bond
unsaturated compound
C(n)H(2n)

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43
Q

benzenes

A
  • unsaturated hydrocarbon
  • perfectly hexagonal, flat ring structure
  • alternating 3 double and single bonds
  • rapid equilibrium model
  • cloud of delocalised e- above/below ring
  • colourless liquid at room temp
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44
Q

formula for cycloalkanes

A

C(n)H(2n)

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45
Q

formula for cycloalkenes

A

C(n)H(2n-2)

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46
Q

saturated hydrocarbons

A

consisting of only single bonds (bonded to the maximum number of hydrogens)

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47
Q

unsaturated hydrocarbons

A

consisting of at least one double or triple bond, thus being more reactive (double bond broken to accommodate addition of more hydrogen atoms)

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48
Q

isomerism

A

same molecular formula but a different spatial arrangement (structural or geometrical)

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49
Q

structural isomerism

A

varying physical structure (can be chain, positional, functional group)

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50
Q

chain isomerism

A

consisting of branches and additional substituents

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51
Q

substituents

A

additional “branches” or “stems”, named as according to where they are and how many carbons they have

52
Q

positional isomerism

A

location of double bonds

53
Q

functional group isomerism

A

location of functional groups

54
Q

geometrical isomerism (cis-trans)

A

structural rotations around C-C double bonds

55
Q

sigma bonds

A
  • single bonds
  • strongest out of all bonds
  • rotational symmetry
    (NO geometrical isomerism)
56
Q

pi bonds

A
  • double or triple bonds
  • weaker and more reactive
    (rotation causes weaker bond to “snap” or break; has different symmetry; geometric isomerism)

much more electron deficient ⇒ much more susceptible to reactions

57
Q

reactivity of hydrocarbons

A

unsaturated hydrocarbons are MORE reactive than saturated hydrocarbons;

for alkanes to react, highly stable sigma bonds must be broken (requiring large amounts of energy to break the stronger bonds) ⇒ large input energy ⇒ relatively higher activation energy ⇒ lower rate of reaction ⇒ lower reactivity

58
Q

substitution reactions

A

type of chemical reaction where an atom or functional group of a molecule is replaced by another atom or functional group

59
Q

substitution of alkanes (energy)

A

UV light supplies sufficient energy to meet the activation energy of the reaction => breaking bonds between halogens (diatomic molecules) => atoms in high energy state can react with alkane to substitute with hydrogen

60
Q

substitution of benzenes (catalyst)

A

halogens involved in substitution aren’t electrophilic enough to break the aromatic nature of benzene, so a catalyst is used to attract electrons for substitution

e.g. nickel catalyst for hydrogenation

61
Q

addition reaction

A

type of chemical reaction where atoms are added across the double bond of an alkene to form an alkane or a substituted alkane, breaking pi bond to form another sigma bond

62
Q

why do addition reactions happen so quickly

A

sigma is stronger and more energetically favourable to form, being more stable, so weaker pi bond will more readily break to form a stabler sigma bond

63
Q

test for unsaturation

A

add few drops of bromine water (orange) or iodine water (brown) into the tested liquid, and shake

two layers would form, since nonpolar hydrocarbons are insoluble in the polar solution (halogens bottom layer)

if unsaturated, would undergo addition reaction and decolorise

64
Q

complete combustion reactions

A

hydrocarbon + excess oxygen → carbon dioxide + water + energy

65
Q

incomplete combustion reactions

A

hydrocarbon + insufficient oxygen → carbon monoxide + water + energy

hydrocarbon + insufficient oxygen → carbon + water + energy

less exothermic than complete combustion, i.e. releasing lots of energy per unit; soot is air pollution, carbon monoxide is poisonous gas

66
Q

pressure

A

force exerted per unit area by collisions between gas particles and interior surfaces of container (gases consist of particles in constant random motion), measured in pascals (Pa)

67
Q

temperature

A

measure of average molecular kinetic energy

68
Q

activation energy

A

minimum amount of energy required for particles to react

minimum amount of energy required to form the activation complex

(energy needed for the bonds of the reactants to break)

69
Q

transition state (activation complex)

A

state corresponding to the highest potential energy (particular configuration) along the reaction coordinate

activation complex = temporary, highly unstable formation (bonds of reactants broken, bonds of products not yet formed)

70
Q

collision theory

A
  • reactants must collide
  • with sufficient, activation energy
  • at appropriate orientation
71
Q

rate of reaction depends on

A
  • frequency of collisions
  • proportion of successful collisions
72
Q

rate of reaction

A

NUMBER OF SUCCESSFUL COLLISIONS OVER A GIVEN TIME

change in concentration of reactants or products per unit time (speed of reaction)

measured using indicators of reactions (mass, colour, volume, pH, concentration) => volume of gas produced, loss of mass, change in transparency

slows over time (concentration of reactants decreases, reduced frequency of collisions)

73
Q

key factors for rate of reaction

A

temperature
surface area
volume / pressure
presence of catalysts
concentration

74
Q

concentration of dissolved reactants

A

higher concentration ⇒ more particles in the same amount of space ⇒ increased frequency of collisions, same proportion of collisions that are successful ⇒ more successful collisions (same chance but more collisions)

75
Q

pressure of gaseous reactants (volume of container)

A

decreased volume, increased pressure ⇒ smaller space for gas particles, closer together ⇒ increased frequency of collisions, same proportion of collisions that are successful ⇒ more successful collisions (same chance but more collisions)

76
Q

surface area of solid reactants

A

(all reactions with solids occur on the surface) split into several pieces ⇒ increased surface area ⇒ increased area for reactants particles to collide with ⇒ increased frequency of collisions, same proportion of collisions that are successful ⇒ more successful collisions (same chance but more collisions)

77
Q

temperature

A

higher temperature ⇒ higher average kinetic energy of particles

increased average velocity of particles ⇒ increased frequency of collisions ⇒ more successful collisions (same chance but more collisions)

colliding with more energy (increased number of collisions having sufficient activation energy) ⇒ increased proportion of collisions that are successful

78
Q

catalyst

A

a substance that speeds up a chemical reaction without being consumed

79
Q

presence of a catalyst

A

offers an alternate pathway for the reaction with a lower activation energy, without directly lowering the activation energy

greater proportion of collisions have sufficient energy to overcome activation energy and form transition complex ⇒ greater proportion of successful collisions

80
Q

maxwell-boltzmann distribution curve

A

demonstrates the proportion of particles that are expected to successfully undergo a collision and thus a chemical reaction

81
Q

why is incomplete combustion is undesirable when using hydrocarbons as fuels

A

less exothermic than complete combustion, i.e. releasing less of energy per unit

soot is air pollution, carbon monoxide is poisonous gas => soot damages engine components overtime and can cause clogs in various pipes etc, direct contributor to global dimming

82
Q

benefits and drawbacks of using hydrocarbons as fuels

A

✅ advantages
- highly calorific / producing large amounts of energy
- readily and cheaply available
- benefitting the economy (gas and oil industry), employment

❌ disadvantages
- non-renewable
- becoming expensive to extract
- negative environmental impact
- enhanced greenhouse effect

    - melting glaciers / increasing water levels / flooding
    - destruction of habitat and extinction of species
    - extreme weather events (earthquakes, heatwaves, bushfires)
- air pollution
83
Q

carbon neutral

A

activity that produces no overall CO2 emissions, by either:

  • reducing emissions (greener method)
  • carbon offsetting (balancing activities that emit CO2 with activities that absorb it e.g. planting trees)
84
Q

biofuels

A

renewable fuels from plant material e.g. agricultural crops

e.g. bioethanol, biodiesel, biogas

85
Q

bioethanol

A

using alcohol from natural fermentation of carbohydrates

86
Q

biodiesel

A

using chemically reacting vegetable oils / animal fats with alcohol

87
Q

biogas

A

gas produced by the breakdown of organic matter in the absence of oxygen

88
Q

advantanges of biofuels (generic)

A
  • burning biofuels is carbon neutral (carbon released = carbon absorbed when plants were growing)
  • renewable processes
  • low storage, transport, distribution costs
  • by-products (pressed seedcake) can be reused and burnt in power stations
89
Q

disadvantanges of biofuels (generic)

A
  • production of biofuels requires energy
  • only few local producers in small quantities; must import
  • high demand for land to plant biofuel crops, resulting in deforestation (habitat loss ⇒ endangerment/extinction)
  • expensive to produce
90
Q

disadvantage of biogas

A

contains some gases as impurities → corrosive to metal parts of internal combustion engines, lower yields due to diluted condition

91
Q

disadvantage of biodiesel

A

not suitable in cold temperatures (solidified)

92
Q

kinetic theory

A

model used to explain the macroscopic behaviour of gases (common physical properties) using understanding of the molecular behaviour of particles [linking macroscopic and microscopic properties]

93
Q

properties of ideal gas

A

hypothetical gas whose molecules occupy negligible volume, thus having no interactions with other gas particles and consequently obeying all ideal gas laws (pV=nRT)

94
Q

ideal gas vs real gas

A

in real gases, there are intermolecular forces (if temperature is lowered or gases are compressed sufficiently then these attractive intermolecular forces will increase and gases will condense to form liquids)

in real gases, particles have some volume, so at zero kelvin the volume will never really be 0 as particles also occupy some volume (whereas in ideal gases, according to assumption, at zero kelvin, volume will be 0)

similarities:
- in rapid continuous random motion
- have kinetic energy
- average kinetic energy proportional to temperature

95
Q

kinetic theory of gases

A
  1. gases consist of identical molecules in rapid, continuous, random motion
  2. attractive and repulsive forces between particles are negligible
  3. particles themselves have negligible volume (average distance between gas molecules is very large)
  4. particles of gas have kinetic energy [Ek=1/2mv^2]
  5. average kinetic energy of gas particles is proportional to temperature and same for all gases (at a specific temperature, maxwell-boltzmann curve will look same)
  6. all collisions are perfectly elastic between gas molecules (no net energy loss)
  7. temperature is a measure of average kinetic energy of particles
96
Q

[explain property using kinetic theory]

takes the shape of its container

A

particles in rapid, continuous, random motion with negligible attractive or repulsive forces ⇒ spread out as far as possible ⇒ occupy entire volume and shape of container

97
Q

[explain property using kinetic theory]

low density

A

particles in rapid, continuous, random motion with negligible attractive or repulsive forces and large amounts of intermolecular space (particle volume negligible) ⇒ spread out as far as possible ⇒ occupy large volumes (maximum volume of container ⇒ distributing mass over a large volume ⇒ low density

98
Q

[explain property using kinetic theory]

pressure

A

gaseous particles in rapid, continuous, random motion with no attraction ⇒ moving in straight lines and colliding with interior surfaces (container walls) ⇒ gas exerts a force per unit area ⇒ pressure

99
Q

[explain property using kinetic theory]

diffusion (readily diffuses through other gases)

A

particles in rapid, continuous, random motion with negligible forces of attraction ⇒ widely spaced ⇒ can easily spread between molecules of any other gas ⇒ can move into the volume between pre-existing gas particles without any attraction between them

100
Q

[explain property using kinetic theory]

compressibility

A

particle volume negligible with negligible repulsive forces ⇒ gas particles spread over larger volume, with large amounts of intermolecular space ⇒ can be compressed into a smaller volume

101
Q

absolute zero

A

0K or 273.15 degrees Celsius, where gas particles stop moving (lowest temperature ⇒ no kinetic energy ⇒ no motion)

102
Q

boyles law

A

at a given constant temperature, the volume is inversely proportional to pressure

103
Q

justify boyles law

A

decreased volume⇒ less room for particles to move around, particles travel less before collisions occur ⇒ increased collision per unit time ⇒ increased collision frequency ⇒ more momentum transferred per unit time to container walls ⇒ force supplied increases ⇒ increased pressure

104
Q

charles law

A

at a given constant pressure, the volumeis directly proportional to temperature

105
Q

justify charles law

A

increased temperature ⇒ increased average kinetic energy of particles ⇒ increased average velocity of particles ⇒ more collisions (increased frequency) but constant pressure ⇒ increased volume (faster particles travel for a longer time to reach container walls)

106
Q

pressure law

A

at a given constant volume, the pressure is directly proportional to temperature

107
Q

justify pressure law

A

increased temperature ⇒ increased average kinetic energy of particles ⇒ increased average velocity of particles ⇒ greater impulse and greater force when colliding with walls due to higher speeds AND more frequent collisions due to greater motion⇒ increased rate and force of particles collisions ⇒ increased force applied per unit area of container walls ⇒ increased pressure

108
Q

combined law

A

p1v1/t1 = p2v2/t2

109
Q

energy

A

the capacity to do work, measured in Joules (J)

110
Q

law of conservation of energy

A

energy can neither be created nor destroyed, only transferred

111
Q

internal energy

A

sum of molecular energies of a system, equal to the sum of kinetic energy due to a molecule’s random motion and its potential energy due to chemical intermolecular and intramolecular bonds

112
Q

thermal energy

A

component of internal energy due to temperature

113
Q

heat

A

transfer of thermal energy from hotter to colder object, until a thermal equilibrium is reached, due to solely temperature difference

114
Q

thermal equilibrium

A

bodies at same temperature with no net flow of thermal energy

115
Q

enthalpy

A

quantitative measure of energy stored within a system, including potential energy (chemical, gravitational, electrical etc) of substance and kinetic energy of its particles

116
Q

enthalpy change of a system (delta H)

A

heat energy exchange with its surroundings at a constant pressure, measured in kJ/mol

delta H = reactants - products

117
Q

exothermic

A

releasing energy from system to surroundings as heat/light (increase in temp)

118
Q

endothermic

A

absorbing energy from surroundings to system as heat (decrease in temp)

119
Q

steps of enthalpy changes

A
  1. bonds breaking (endothermic process ⇒ increases enthalpy)when bonds break, work needs to be done as a force needs to be applied to overcome the electrostatic forces of attraction between electrons and protons in the bonds, so energy is required; stronger bonds require larger force to be applied → larger work (J) → more energy required
  2. activation complex / transition state (temporary, highly unstable molecule)
  3. bond formation (exothermic process ⇒ decreases enthalpy)products (bonded atoms) are such that resultant molecule is stable (complete valence shell achieved), thus containing less potential energy, and thus an increase in kinetic energyalso represents the reverse reaction corresponding to breaking a bond (which requires heat), and thus releases heat

bond breaking and bond making represent energy consumption and energy production respectively ⇒ ALL REACTIONS have a net change in CHEMICAL SYSTEM’S TOTAL ENERGY

120
Q

bond energy

A

amount of energy required to break a bond OR amount of energy released in formation of bond, corresponding to the stability of the bond, and thus its strength

121
Q

thermochemical equations

A

representing enthalpy changes in chemical reactions

122
Q

melting

A

the endothermic physical process in which heat flows into a solid system to convert the state of the system into liquid, only occurring when the solid reaches its melting point temperature

123
Q

boiling

A

the endothermic physical process in which heat flows into a liquid system to convert the state of the system into gas, only occurring when the liquid reaches its boiling point temperature

124
Q

condensation

A

the exothermic physical process in which heat flows OUT of a gaseous system to convert the state of the system into liquid, only occurring when the gas is cooled to lower than its boiling point temperature

125
Q

freezing

A

the exothermic physical process in which heat flows OUT of a liquid system to convert the state of the system into solid, only occurring when the liquid is cooled to lower than its melting/freezing point temperature

126
Q

sublimination

A

the endothermic physical process in which heat flows INTO a solid system to convert the state of the system into GAS, skipping the intermediate liquid phase due to the extremely low pressure of the system

AECHE (8 decks)

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