mocks Flashcards
(97 cards)
1
Q
what are the features of a hydrogen emission spectrum
A
- discrete lines (line spectrum) = energy levels are discrete
- lines converge toward high frequency = energy levels are getting closer at high energy
2
Q
what are the names of the series in an electron transition from higher to lower
A
- Lyman series: N = 1
- Balmer series: N = 2
- Paschen series: N = 3
3
Q
how to calculate ionisation energy (in kJ mol^-1)
A
c = v * λ
E = h * v
v is frequency
λ is convergence limit (the frequency at which the spectral lines converge)
c = speed of light
h = Planck’s constant
E = energy
4
Q
what are the rules of electron configuration
A
- Aufbau principle - in the ground state of an atom or ion, electrons first fill subshells of the lowest available energy, then fill subshells of higher energy
- Pauli exclusion principle - arrows must be pointing in diff directions
- Hund’s Law - electrons always enter an empty orbital before they pair up
–> i.e. chromium (Ar 4s1 3d5) and copper (Ar 4s1 3d10)
5
Q
sig fig and d.p.
A
if +/-, then dp = least dp in question
if x/➗, then sf = least sf in question
6
Q
absolute uncertainty formula
A
(max value - min value)/2
7
Q
percentage uncertainty formula
A
absolute uncertainty/total measured value
8
Q
what is absolute error
A
= absolute uncertainty
9
Q
percentage error formula
A
|expected value - actual value|/actual value * 100
10
Q
avogadro’s constant
A
at constant temp and pressure, volume is directly proportional to moles
(volume ratio = mole ratio)
11
Q
SI units of P, V, n, T
A
P = kPa or Pa (kPa = 1000 Pa)
V = dm^3 or m^3 (m^3 = 1000dm^3)
n = mol
T = K
12
Q
properties of ideal gases
A
negligible volume
no intermolecular force
have elastic collision
13
Q
mono vs diprotic acids
A
basically number of H
14
Q
lowest energy transition on visible spectrum
A
n=3 to n=2
15
Q
Which region of the electromagnetic spectrum is used to identify hydrogen environments in a molecule?
A
radio waves
16
Q
Distinguish ultraviolet light from visible light in terms of wavelength and energy
A
shorter wavelength, higher energy
17
Q
What is the maximum number of electrons that can occupy the 4th main energy level in an atom
A
32
18
Q
what is the expression to find the maximum number of electrons that can occupy the nth main energy level in an atom
A
2n^2
19
Q
factors affect IE
A
- nuclear charge
- number of energy levels
- atomic radius
20
Q
trends of first IE
A
across a period:
- nuclear charge increases
- atomic radius decreases
down a group
- IE decreases
21
Q
methods to determine rate of reaction
A
change in volume of gas
change in mass
change in colour
change in pH
change in electrical conductivity
titration
22
Q
factors affecting equilibrium position
A
- concentration = shifts away from increase in conc.
- pressure = shifts to side with less moles
- temperature = increase in temp shifts to endo.
23
Q
molar gas volume formula
A
gas volume = moles * 22.7
24
Q
percentage yield formula
A
actual/theoretical * 100
25
atom economy formula
(Mr of desired product*coefficient) /(sum of Mr for all reactants*coefficients) * 100
26
equilibrium constant formula
K = ([C]^c[D]^d)/([A]^a[B]^b)
27
what does Kc tell us?
Kc >>> 1 = reaction almost complete
Kc <<< 1 = reaction barely occurs
28
what factors affect Kc?
only temperature
ENDO: T↑, Kc↓
EXO: T↓, Kc↑
29
define ionic bond
ELECTROSTATIC ATTRACTION between positive and negative ions
30
define covalent bond
ELECTROSTATIC ATTRACTION between shared pair of electrons and both nuclei of the bonding atoms
31
what is a coordination bond/dative covalent bond
each atom donates 1 e- = regular
both e- from same atom = coordination bond
both have the same bond length and strength
32
vsepr 5 e- domains on central atom
trigonal bypyramidal
33
types of intermolecular forces
london dispersion (all molecules)
dipole dipole (polar molecules)
hydrogen bond (N, O, F with H)
34
physical properties of simple molecules
1. relatively low B.P./high volatility (i.e. easy to evaporate
- due to weak IMF between molecules that require little energy to break
- stronger IMF = higher B.P.
2. solubility
- non-polar molecules dissolve in non-polar solvents
- polar molecules dissolve in polar solvents
- H-bond molecules dissolve in water
3. poor electrical conductivity = no mobile ions/delocalised electrons
35
diamond
each c atom covalently bonded to 4 others
- non conductor
- very efficient thermal conductor
- brittle
- high melting point
- used in jewellery, tools and machinery
36
graphite
each c atom covalently bonded to 3 others, forms sheets
- conductor
- not a good thermal conductor
- brittle
- very high melting point (most stable allotrope of carbon)
- used in pencils + electrolysis
37
graphene
each c atom covalently bonded to 3 others, forms ONE sheet
- VERY good conductor
- BEST thermal conductor
- very flexible and very strong
- very high melting point
- used in touch screens and high performance electronic devices
38
C60 fullerene
sphere made of 60 carbons
- poor conductor
- very low thermal conductivity
- in the form of powder, very light and strong
- low melting point (most stable allotrope of carbon)
- used in carbon nanotubes, catalysts
39
average bond enthalpy
amount of energy required to break one mole of covalent bonds in gaseous molecules
40
atomic radius trends
decrease across period (similar shielding, protons inc.)
increase down group (shells inc., shielding inc.)
41
ionic radius trends
increases down the group (shielding inc.)
42
ionisation energy trends
increase across period (shielding similar, proton number inc.)
decrease down group (atomic radius inc.)
43
electron affinity trends
increase across period (proton number inc., shielding similar)
decrease down group (more shielding, more shells)
44
electronegativity trends
increase across period (proton number inc. similar shielding)
decrease down group (more shielding, more shells)
45
melting and boiling point trends
depends on bonding and structure
- giant covalent highest
- compare ionic radius, charge and number of delocalised e-
group 1: decrease, ionic radius inc., charge is the same
group 17: intermolecular forces, LDF is the strongest in iodine
46
reaction between group 1 alkali and water
metal + water --> metal hydroxide + water
47
reactivity trends of alkali metals
increase down group
--> more shells, more shielding, tendency to lose electron is greater
--> atomic radius increases
48
reaction between group 17 halogens and halides
displacement reaction
49
methods of reducing environmental impact of so2 and so3
pre combustion = remove sulfure from fossil fuels
post combustion = remove so3 and so3
50
methods of reducing environmental impact of no2
post combustion = catalytic converter to remove no2
51
what is a homologous series
same general formula, same functional group, similar chemical properties, gradual change in physical properties, successive members differ by a CH2 group
52
structural isomers
same molecular formula but diff structural formula
53
chain isomers
difference in main chain and branching
54
position isomers
same functional group but different positions
55
functional group isomers
different functional groups
56
classification of alcohols/halogenoalkanes/nitrogen atoms in amines
primary = 1 alkyl group/1º (i.e. CH3, C2H2)
secondary = 2º
tertiary = 3º
57
melting points of cis and trans isomers
cis: higher boiling point due to net dipole movement
trans: higher melting point due to more regular packing
58
chiral carbon
carbon bonds to 4 different atoms/groups of atoms
59
optical isomers
basically mirror image of the same molecule, have the same physical properties except optical activity
60
how to identify optical isomers?
polarimeter
61
bronsted-lowry acids and bases
acids = proton donor
bases = proton acceptors
62
conjugate acid base pairs
CH3COOH (acid) and CH3COO- (base)
63
amphoteric
can react w/ both acid and base (i.e. Al2O3)
64
amphiprotic
can act as both a bronsted lowry acid or base
65
pKw
pKw = 14, pKw = pH + pOH
66
pH of strong acids and bases
strong = completely dissociates
67
sig figs and d.p. in pH calculations
number of sig figs in conc = number of d.p. in pH
68
strengths of acids and bases and how does this relate to being proton donors/acceptors and the type of conjugate bases produced?
strong acids =
✅ proton donor, ❌ conjugate base
weak acids =
❌ proton donor, ✅ conjugate base
strong bases =
✅ proton donor, ❌ conjugate acid
weak bases =
❌ proton donor, ✅ conjugate acid
69
how to differentiate between strong and weak acids
1. measure pH using a pH meter
2. measure electrical conductivity
3. react by metal (or metal carbonate)
70
acid + metal
salt + hydrogen
71
acid + metal oxide/hydroxide
salt + water
72
acid + metal carbonate
salt + water + carbon dioxide
73
acid + ammonia
ammonium salts
74
what are oxidation and reduction in reference to oxidation states?
oxidation = ++++ oxidation state
reduction = ----- oxidation state
75
redox half equations in neutral condition
e- on the left = reduction
e- on the right = oxidation
write half equations as normal
76
redox half equations in acidic conditions
oxidation as normal
reduction = add H+ to left and H2O on right
77
common oxidising agents during oxidation of alcohol
acidified potassium dichromate (VI)
- K2Cr2O7
acidified potassium manganate (VII)
- KMnO411
78
oxidation of alcohol process
PRIMARY
alcohol --> partial oxidation using distillation: aldehyde (H-C=O) --> complete oxidation by reflux carboxylic acid (COOH)
SECONDARY
alcohol --> ketone (-CO-)
reflux
TERTIARY = N/A
78
stages of nucleophilic substitution
78
stages of radical substitution
alkane + halogen --> halogenoalkane + hydrogenohalide
must be in the presence of UV light
example: CH₄ + Br₂ → CH₃Br + HBr
initiation
1. Br-Br → Br· + Br·
curly fish hook
propagation
2. Br· + CH₄ → CH₃· + HBr
3. CH₃· + Br₂ → CH₃Br + Br·
termination (when all radicals are used up)
1. Br· + Br· → Br₂
2. CH₃· + CH₃· → C₂H₆
3. Br· + CH₃·
79
rate equation
aA + bB → cC + dD
Rate = k (rate constant) [A]^m [B]^n
m and n = order
80
stages of electrophilic substitution
81
rate graphs
order = 0
rate constant with conc. of substance
rate inversely proportional with time
order = 1
rate directly proportional with conc. of substance
rate like a inverse sagging curve against time
order = 2
rate is directly proportional with [A]^2
rate is proportional, sagging curve upwards against [A]
rate like a inverse sagging curve against time, but steeper than first order
82
formal charge formula (for a specific atom in a molecule)
no. of valence in the atom - 1/2 (no of bonding e-) - non bonding e-
83
sigma bond and pi bond
sigma =head on combo of atomic orbitals where the electron density is concentrated along the bond axis
pi = formed by the lateral combo of p-orbitals where the e- density is concentrated on opposite sides of the bond axis
sigma = all single bonds
all double bonds = first bond is sigma, then all are pi
84
hybridisation
2 e- domains, linear domain geometry = sp (2 and 2)
3 e- domains, trigonal planar domain geometry = sp (3 and 1)
4 e- domains, tetrahedral domain geometry = sp (4 and 0)
85
calculating ΔH using ΔHf
products - reactants
86
calculating ΔH using ΔHc
reactants - products
87
born haber cycle order
1. ΔHf
2. ΔHatom (s->g, diatomic -> 2monoatomic)
3. IE (positive ion)
4. Electron affinity (negative ion)
5. Lattice enthalpy
88
methods of calculating calorimetry
1. Q = mcΔT
2. Hess's Law (ΔHf P-R)
3. bond enthalpy (R-P)
4. lattice enthalpy
89
ΔS of a system
products - reactants
90
ΔS of surroundings
= - ΔH/T
if + = exo
if - = endo
91
what does ΔG = ΔH - TΔS tell us
ΔG is in J
1. ΔH +, ΔS + spontaneous at high temp
2. ΔH +, ΔS - non-spontaneous at any temp
3. ΔH -, ΔS + spontaneous at any temp
4. ΔH -, ΔS - spontaneous at low temp
92
reaction quotient (Qc)
find conc. at any time (kc only at equilibrium)
same formula as kc
92
what does qc tell us?
Qc > Kc = reaction shifts to right
Qc < Kc = reaction shifts to left
Qc = Kc = reaction is at equilibrium
93
bonding triangle
Δelectronegativity = Y-axis
Average electronegativity = X-axis
94
draw a pH curve
see answer on goodnotes
