sem 2 exam (unit 2 content) Flashcards
1
Q
acids
A
substances that ionise to release H+ ions when dissolved in water
2
Q
bases
A
substances that produce OH- ions in solutions
3
Q
properties of acids:
pH, solutions, taste, feel
A
pH < 7, forms H+ ions, sour, stinging/burning skins
4
Q
properties of bases:
pH, solutions, taste, feel
A
pH > 7, forms OH- ions, bitter, slippery feel
5
Q
indicators
A
chemicals that change colour in the presence of an acid and alkali
6
Q
indicating acids
A
blue litmus paper becomes red, OR
red/orange with universal indicator
7
Q
indicating bases
A
red litmus paper becomes blue, OR
blue/purple with universal indicator
8
Q
ionisation
A
where a covalent molecule splits into ions (i.e. all acids)
9
Q
dissociation
A
where an ionic compound dissolves in water
10
Q
strong vs weak
A
strong: completely ionise/dissociate in water
weak: only partially ionise/dissociate
11
Q
HCl
A
strong acid
12
Q
HNO3
A
strong acid
13
Q
H2SO4
A
strong acid
14
Q
CH3COOH
A
weak acid
15
Q
H3PO4
A
weak acid
16
Q
H2CO3
A
weak acid
17
Q
common strong bases
A
group 1 and 2 oxides and hydroxides
18
Q
NH3
A
weak base
19
Q
Na2CO3
A
weak base
20
Q
NaHCO3
A
weak base
21
Q
concentrated
A
high molarity
22
Q
diluted
A
low molarity
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)
24
Q
pH scale
A
inverse logarithmic measure of hydrogen ion concentration
25
acid + metal
salt + hydrogen gas
26
acid + metal oxide
salt + water
27
acid + metal hydroxide
salt + water
28
acid + metal carbonate
salt + water + carbon dioxide
29
acid + metal hydrogencarbonate
salt + water + carbon dioxide
30
base + ammonium salt
salt + water + ammonia (pungent gas)
31
base + nonmetal oxide
salt + water
32
self-ionisation (autoionisation) of water
equilibrium reactions that water undergoes on its own, by ionising and recombining constantly
33
ionic product of water (Kw)
equillibrium constant for the self-ionisation of water
Kw = [H+]*[OH-] = 1 x 10^-14 at 25deg
34
neutrality when
concentration of hydrogen = hydroxide
35
why is carbon able to form such a diverse range of compounds?
2, 4 electron configuration (4 valence) ⇒ ability to covalently form 4 more single bonds ⇒ variety of arrangements (single, double, triple) ⇒ forms many compounds
36
hydrocarbons
organic compounds, based on hydrogen and carbon atoms
37
aliphatic
carbon atoms arranged in open chains ⇒ can be saturated or unsaturated
38
aromatic
carbon atoms arranged in closed rings
39
functional groups
atom or group of atoms responsible for typical chemical reactions, determining patterns of reactivity
40
homologous series
group of molecules with same functional group, but different number of carbons in the main chain
41
alkanes
single bonds only
saturated compounds
C(n)H(2n+2)
42
alkenes
at least one double bond
unsaturated compound
C(n)H(2n)
43
benzenes
- 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
44
formula for cycloalkanes
C(n)H(2n)
45
formula for cycloalkenes
C(n)H(2n-2)
46
saturated hydrocarbons
consisting of only single bonds (bonded to the maximum number of hydrogens)
47
unsaturated hydrocarbons
consisting of at least one double or triple bond, thus being more reactive (double bond broken to accommodate addition of more hydrogen atoms)
48
isomerism
same molecular formula but a different spatial arrangement (structural or geometrical)
49
structural isomerism
varying physical structure (can be chain, positional, functional group)
50
chain isomerism
consisting of branches and additional substituents
51
substituents
additional “branches” or “stems”, named as according to where they are and how many carbons they have
52
positional isomerism
location of double bonds
53
functional group isomerism
location of functional groups
54
geometrical isomerism (cis-trans)
structural rotations around C-C double bonds
55
sigma bonds
- single bonds
- strongest out of all bonds
- rotational symmetry
(NO geometrical isomerism)
56
pi bonds
- 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
reactivity of hydrocarbons
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
substitution reactions
type of chemical reaction where an atom or functional group of a molecule is replaced by another atom or functional group
59
substitution of alkanes (energy)
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
substitution of benzenes (catalyst)
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
addition reaction
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
why do addition reactions happen so quickly
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
test for unsaturation
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
complete combustion reactions
hydrocarbon + excess oxygen → carbon dioxide + water + energy
65
incomplete combustion reactions
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
pressure
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
temperature
measure of average molecular kinetic energy
68
activation energy
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
transition state (activation complex)
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
collision theory
- reactants must collide
- with sufficient, activation energy
- at appropriate orientation
71
rate of reaction depends on
- frequency of collisions
- proportion of successful collisions
72
rate of reaction
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
key factors for rate of reaction
temperature
surface area
volume / pressure
presence of catalysts
concentration
74
concentration of dissolved reactants
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
pressure of gaseous reactants (volume of container)
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
surface area of solid reactants
(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
temperature
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
catalyst
a substance that speeds up a chemical reaction without being consumed
79
presence of a catalyst
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
maxwell-boltzmann distribution curve
demonstrates the proportion of particles that are expected to successfully undergo a collision and thus a chemical reaction
81
why is incomplete combustion is undesirable when using hydrocarbons as fuels
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
benefits and drawbacks of using hydrocarbons as fuels
✅ 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
carbon neutral
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
biofuels
renewable fuels from plant material e.g. agricultural crops
e.g. bioethanol, biodiesel, biogas
85
bioethanol
using alcohol from natural fermentation of carbohydrates
86
biodiesel
using chemically reacting vegetable oils / animal fats with alcohol
87
biogas
gas produced by the breakdown of organic matter in the absence of oxygen
88
advantanges of biofuels (generic)
- 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
disadvantanges of biofuels (generic)
- 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
disadvantage of biogas
contains some gases as impurities → corrosive to metal parts of internal combustion engines, lower yields due to diluted condition
91
disadvantage of biodiesel
not suitable in cold temperatures (solidified)
92
kinetic theory
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
properties of ideal gas
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
ideal gas vs real gas
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
kinetic theory of gases
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
[explain property using kinetic theory]
takes the shape of its container
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
[explain property using kinetic theory]
low density
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
[explain property using kinetic theory]
pressure
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
[explain property using kinetic theory]
diffusion (readily diffuses through other gases)
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
[explain property using kinetic theory]
compressibility
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
absolute zero
0K or 273.15 degrees Celsius, where gas particles stop moving (lowest temperature ⇒ no kinetic energy ⇒ no motion)
102
boyles law
at a given constant temperature, the volume is inversely proportional to pressure
103
justify boyles law
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
charles law
at a given constant pressure, the volumeis directly proportional to temperature
105
justify charles law
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
pressure law
at a given constant volume, the pressure is directly proportional to temperature
107
justify pressure law
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
combined law
p1v1/t1 = p2v2/t2
109
energy
the capacity to do work, measured in Joules (J)
110
law of conservation of energy
energy can neither be created nor destroyed, only transferred
111
internal energy
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
thermal energy
component of internal energy due to temperature
113
heat
transfer of thermal energy from hotter to colder object, until a thermal equilibrium is reached, due to solely temperature difference
114
thermal equilibrium
bodies at same temperature with no net flow of thermal energy
115
enthalpy
quantitative measure of energy stored within a system, including potential energy (chemical, gravitational, electrical etc) of substance and kinetic energy of its particles
116
enthalpy change of a system (delta H)
heat energy exchange with its surroundings at a constant pressure, measured in kJ/mol
delta H = reactants - products
117
exothermic
releasing energy from system to surroundings as heat/light (increase in temp)
118
endothermic
absorbing energy from surroundings to system as heat (decrease in temp)
119
steps of enthalpy changes
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 energy*
*also 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
bond energy
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
thermochemical equations
representing enthalpy changes in chemical reactions
122
melting
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
boiling
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
condensation
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
freezing
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
sublimination
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
