1
Q
law of conservation of mass
A
products have the same mass as starting materials, energy is transformed, not created or destroyed
2
Q
atomic theory
A
all matter is composed on atoms which cannot be created or destroyed, only rearranged
3
Q
elements
A
primary constituents of matter, which cannot be chemically
broken down into simpler substances
4
Q
compounds
A
consist of atoms of different elements chemically bonded together in a fixed ratio
5
Q
mixture (homogeneous vs heterogeneous)
A
contain 2+ element or compound in no fixed ratio, which are not chemically
bonded and so can be separated by physical methods
homo: uniform composition
hetero: non-uniform
6
Q
separating mixtures: filtration
A
- mixture poured thru paper filter or porous material
- separates liquids from solids
7
Q
separating mixtures: dissolution/solvation
A
mixture added to water or organic solvent
- separates soluble substances from insoluble substances
8
Q
separating mixtures: crystallization
A
- mixture dissolved in hot water/organic solvent then cooled, forming crystals -> to be isolated thru filtration
- removes more soluble substances from less soluble substances (crystals)
9
Q
separating mixtures: evaporation/distillation
A
- mixture heated until 1+ components vaporize
- removes volatile liquids from solids/less volatile liquids
10
Q
separating mixtures: paper chromatography
A
- mixture placed on piece of paper -> one side is submerged in water/solvent moving components along
- more soluble components move further, less soluble components move less
11
Q
characteristics of a solid
A
- fixed volume
- fixed shape
- cannot be compressed
- strong attractive forces
- particles vibrate
12
Q
characteristics of a liquid
A
- fixed volume
- no fixed shape
- cannot be compressed
- medium attractive forces
- particles vibrate, rotate, move around a little
13
Q
characteristics of a gas
A
- no fixed volume
- no fixed shape
- can be compressed
- little attractive forces
- particles vibrate, rotate, move around lots
14
Q
solid -> gas
A
sublimation
15
Q
endothermic vs exothermic changes of state
A
endo: when substance changes from more condensed to less condensed state (ex. L->g)
exo: less condensed -> more condensed
16
Q
gas -> solid
A
deposition
17
Q
what is temperature
A
measure of average kinetic energy of particles
18
Q
why is there no temperature change during state changes
A
added energy is used to disrupt solid lattice -> overcome intermolecular forces bw molecules
19
Q
characteristics of a nucleus
A
- very small compared to atom
- highly dense structure, contains all mass of atom (e- surround the nucleus, mass is negligible)
- has positive charge
20
Q
elementary charge
A
electric charge carried by a single electron
21
Q
nuclear symbol notation, how to find # protons/neutrons
A
top number (A) reps mass number, bottom number (Z) reps atomic number
protons: Z
neutrons: A-Z
22
Q
isotopes
A
atoms of the same element with different numbers of neutrons
- diff mass number, same atomic number
23
Q
properties of isotopes
A
- similar chemical properties
- diff physical properties (ex. density as one isotope is heavier than the other)
24
Q
natural abundance (NA)
A
percentage of an isotope’s atoms among all atoms of the given element found on our planet
25
heavy vs light particles in mass spectrometer
will be deflected the least vs will be deflected the most
26
how are emission spectra produced
by atoms emitting photons when electrons in excited states return to lower energy levels
27
continuous spectrum
- occurs when light passes thru prism
- consists of all wavelengths, appearing as continuous series of colours w no gaps
- wavelengths from 400-700 (visible light)
28
emission spectrum
- pure gas elements at high voltage, reduced pressure emits light
- this light passes thru prism, producing coloured lines on dark background
29
absorption spectrum
- when cold gas is placed bw prism and source of visible light
- series of dark lines in continuous spectrum
30
colour with highest/lowest frequency in visible colour spectrum
highest: violet -> most energy, shortest wavelength
lowest: red -> least energy, longest wavelength
31
ground state (stationary state), how many electrons can it hold?
- electrons occupying fixed circular orbits around nucleus (n=1)
- do not emit energy
- most stable state
32
how do electrons go from ground state to excited state?
absorbing specific amount of energy thats exactly equal to the difference b/w 2 states
33
what does energy level 1 produce?
UV radiation
34
what does energy level 2 produce?
visible light
35
what does energy level 3+ produce?
infrared radiation/light
36
hydrogen emission spectra
red (656nm), blue-green (486nm), blue-violet (434nm), violet (410nm).
37
how many electrons can each energy level/state hold?
n=1: 2
n=2: 8
n=3: 8
n=4: 32
38
atomic orbital
region in space where there is a high probability of finding an electron
39
s orbital vs p orbital shapes
cluster of dots or sphere enclosing 99% of dots vs dumbbell shaped, 3 of them on the 3 axises (x, y, z)
40
energy levels of atomic orbitals + number of orbitals in each level
s (1) < p (3) < d (5) < f (7)
41
principal quantum number relationship to total number orbitals and maximum number of e-
- principal quantum number (n) can hold n types of orbitals in increasing energy levels
- ex. n=2 holds s, and p orbitals -> 4 total orbitals -> 8 total e- (2 e- per orbital)
42
electron configuration
arrangement of e- in orbitals
43
pauli's exclusion principle
only two e- can occupy the same atomic orbital and those electrons must have opposite spins
- bcus like charges repel, but with opposite spins they behave like magnets
44
degenerate orbitals
orbitals with the same energy (atomic orbitals of the same type in one sublevel have equal energy)
45
hund's rule
every degenerate orbital in a sublevel is singly occupied before any orbital is double occupied + all e- in singly occupied orbitals have the same spin
46
aufbau's principle
as e- are added to atoms, the lowest available energy orbitals fill before higher energy orbitals do
47
order of filling of energy orbitals
1s<2s<2p<3s<3p<4s<3d<4p
48
3 ways to show electron configuration
full electron configuration, condensed electron configuration, orbital filling diagram
49
ionization energy + equation
minimum energy required to eject an e- out of a neutral atom or molecule in ground state
X(g) + energy -> X^+(g) + e-
50
trend of first ionization energy
- IE increases across period as # of protons inc. so outer e- are held closer to nucleus by larger nuclear charge
- decreases down a group as # of sublevels inc. causing more shielding to outer e-
51
what would first ionization energies look like in a graph
group 1 –> troughs
group 18–> peaks
- after every energy level shell, there is a large drop
- a smaller drop occurs when electrons start pairing (ex. b/w N and O) and after 2s orbital fills
- general upwards trend
52
what is the limit of convergence on an emission spectrum
- as principal quantum numbers of energy levels increase, distance bw levels converges into continuum (infinity)
- represents ionization
53
successive ionization + graph
- energy needed to remove second, third etc. e- from an atom
- requires more energy as # of protons exceeds # of e-
- graph increases, large increase when removing e- from next orbital
54
mole
SI unit of an amount of substance
- contains avogadro's number of elementary entities (atom, molecule, ion, electron etc) of that substance
55
relative atomic mass vs relative molecular/formula mass
ratio of mass of certain ATOM to one twelfth of the mass of a carbon 12 vs. ratio of mass of a MOLECULE or other multiatomic species to one twelfth the mass of a carbon 12 atom
56
calculating relative molecular mass of hydrates
hydrate: compound w water molecules forming coordination bonds
- Mr includes mass of compound and the water WITH the coefficient
57
empirical vs. molecular formula
shows simplest ratio of atoms present in substance vs. actual number of atoms of each elements in substance
58
solutions
homogenous mixtures of 2+ compounds
- contains solvent and 1+ solutes
59
molar concentration (molarity)
ratio of amount of solute to volume of solution
60
purpose of a calibration curve
to relate a measurable property (ex. absorbance, pH, conductivity) of solution to concentration of the solute
61
avogadro's law
states that equal volumes of any two gases at the same temp and pressure contain equal numbers of molecules
62
assumptions of the ideal gas model
1. molecules of gas are in constant random motion
2. collisions bw molecules are perfectly elastic (no energy lost)
3. volume of gas molecules is negligible compared to volume of container they occupy
4. no intermolecular forces bw gas particles
5. kinetic energy is directly proportional to kelvin temperature
63
boyle's law
at a constant temp, the pressure of a given amount of gas is inversely proportional to its volume (ex. double pressure halves volume)
64
when do gases deviate from ideal gas behaviour
- low temp: kinetic energy of gas molecules is reduced so as they collide w each other, intermolecular forces form, may not rebound elastically
- high pressure: more molecules in reduced space -> volume of molecules becomes significant part of volume of gas
65
ideal conditions for ideal gas conditions
- low pressure, high temp
- keep molecules far apart, preventing interaction
66
molar volume of an ideal gas
- constant at specified temp and pressure
- at STP, molar volume of an ideal gas is 22.7dm/mol
67
types of chemical bonds and structures
ionic, covalent (molecular covalent, covalent network), metallic
68
properties of ionic compounds
crystalline, brittle, poor electrical conductors when solid, good electrical conductors when molten or dissolved
69
cations vs anions
ions with more protons than electrons (pos. charged) vs. ions with more electrons that protons (neg charged)
70
isoelectronic
two different species with the same electron configuration (typically due to ions ex. Na+ and Ne)
71
how are isoelectronic species created
- resulting cations and anions have noble gas configuration
- cations go thru oxidation and anions go thru reduction
72
the formation of an ionic compound from its elements is a _____ reaction
redox
73
octet rule
main-group elements tend to bond in ways that provide them with eight valence electrons, like a noble gas
74
hydrogen ion vs hydride ion
H+ (when H loses an e-) vs. H- (when H gains an e-)
75
how are ionic bonds formed
by electrostatic attractions between oppositely charged
ions
76
how to estimate if a bond is ionic
- if the difference in electronegativity bw the two elements is large (>1.8)
77
naming ionic compounds
cation + root of anion + ide
78
ionic lattices
- in ionic crystals ions are arranged in a lattice -> rep. as empirical formula
- continuous 3D network of repeating units of positive and negative ions
- each ion attracts all opposite charges species surrounding it
79
lattice enthalpy
- enthalpy change of the formation of gaseous ions from one mole of a solid lattice
- value tells us how strong ionic bonds are in a particular ionic lattice
80
factors affecting lattice enthalpy
- ionic radius: larger ionic radius = decreases electrostatic attraction
- ionic charge: greater ionic charge = increases electrostatic attraction
81
volatility of ionic compounds
- non-volatile
- strong electrostatic forces of attraction bw ions must be overcome
- also high boiling and melting points
82
electrical conductivity of ionic compounds
- substances must have charged particles able to move
- has charged particles, but ions in solid lattice are not mobile
- when molten/aq, cations and anions can move, making it conductive
83
solubility of ionic compounds
- typically soluble in polar solvents (ex. water), insoluble in nonpolar solvents
- water molecules point partial neg charges at cation, partial positive charges at anions -> ions pulled out of lattice, surrounded by water molecules
84
why are some ionic compounds insoluble in water
- when electrostatic attraction s bw cations and anions in lattice are stronger than association bw ions and water molecules (ex. calcium carbonate, silver chloride)
85
formation of covalent bonds
formed by electrostatic attraction bw a shared pair of electrons and the positively charged nuclei.
86
lewis formula rules
1.only valence e- shown
2. e- arranged in pairs
3. each pair of e- shsared bw 2 atoms is a covalent bond
4. e- in bond are positioned in region bw the 2 atoms involved in bond
5. nonbonding e- (lone pairs) positioned away from bonding region
87
bond lengths + strengths of single, double, triple bonds
s: low strength, long
d: medium strength, medium length
t: high strength, short
88
coordination bond
covalent bond in which both the electrons of the shared pair originate from the same atom
89
coordination bonds in transition metal complexes
- coordination bonds to hold together central metal to surrounding atoms/groups of atoms called ligands
- all ligands have a lone pair of e- to coordination bond to metal ion
90
what does VSEPR stand for
valence shell electron pair repulsion
91
what 3 rules is the VSEPR model based on
1. electron pairs repel each other so they are as far apart from each other as possible
2. lone pairs occupy more space than bonding pairs (single bonds)
3. double and triple bonds occupy more space than single bonds
92
electron domain
region of high electron density due to presence of e- pairs (ex. lone pair, single/double/triple bonds)
93
central atom with 2BP/0LP- electron group arrangement + angles
linear, 180 degrees b/w electrons
- 2 electron domains
94
95
central atom with 3BP/0LP- electron group arrangement + angles
trigonal planar, 120 degrees
- 3 electron domains
96
central atom with 2BP/1LP- electron group arrangement + angles
bent, <120 degrees
- lone pair takes up more space cus it exerts a stronger repulsion
97
central atom with 4BP/0LP- electron group arrangement + angles
tetrahedral, 109.5 degrees
98
central atom with 3BP/1LP- electron group arrangement + angles
trigonal pyramidal, <109.5 degrees (107.5)
99
central atom with 2BP/2LP- electron group arrangement + angles
bent/v-shaped, «109.5 degrees (104.5)
100
what does bond polarity arise from
- difference in electronegativities of bonded atoms
101
dipole moment, how to know where it points
- separation of charge bw two non-identical bonded atoms
- arrow starts as plus sign at less electronegative atom, points towards more electronegative atom (or use partial charges)
102
polar covalent vs pure covalent bond
when electronegativities of 2 atoms in a bond are different, causing partial charges vs. when electronegativities are the same, non-polar
103
what does molecular polarity depend on
both bond polarity and molecular geometry
104
molecular covalent vs covalent network (aka giant covalent)
discrete groups of covalently bonded atoms (molecules) vs atoms held together in a continuous 3D lattice
105
allotropes
structure of the same element that have different bonding and structural patterns
- different chemical and physical properties
106
diamond as an allotrope of carbon, structure + properties
- each C atom bonded to 4 other C atoms in tetrahedral arrangement
- poor electrical conductor (no mobile e-), great thermal conductor, durable, high refractive index
107
graphite as an allotrope of carbon, structure + properties
- layers of sheets made from C atoms -> each C is bonded to 3 other C atoms in trigonal planar
- good electrical conductor (one e- per carbon is delocalized), weak forces of attraction bw sheets
108
graphene as an allotrope of carbon, structure + properties
- single sheet off graphite
- good electrical conductor, flexible, lightweight, super strong
109
fullerenes + buckyballs as allotropes of carbon, structure + properties
- fullerenes area groups of C allotropes arranged in hexagonal and pentagonal rings ->includes carbon nanotubes and buckyballs
- BB: C60 -> looks like a soccer ball, low boiling point (weak intermolecular forces)
110
silicon structure + properties
- 3D lattice where each silicon atom is bonded to 4 other silicon atoms in tetrahedral
- high strength, high melting+boiling points (not as high as C), semi-conductor
111
silica (silicon dioxide) structure+properties
- each silicon atom is bonded to 4 Os and each O is bonded to 2 silicons
- crystalline structure: quartz
- hard, insoluble in water, high melting point, poor conductor of electricity and heat as a solid
112
van der Waals
intermolecular forces of LDFs, dipole induced dipole forces, dipole-dipole forces
113
how do london dispersion forces occur
- all molecules experience it
- uneven e- distribution causes temporary instantaneous dipoles in an atom, inducing dipoles in nearby atoms
- partial negatives and partial positives will attract each other
114
2 factors that affect the strength of london dispersion forces
- number of electrons (more e- = stronger LDFs)
- molecular size/mass (larger size= more e-=stronger LDFs)
115
how are london dispersion forces affected by the shape of a molecule
- "sticky sausage"
- longer molecules have a larger area of interaction and better contact with other molecules compared to shorter ones -> stronger LDFs, higher BP and MP
116
how dipole-induces dipole forces occur
- bw a polar molecule and a nearby non-polar molecule
- permanent dipole in polar molecule induces formation of temp. dipole in non-polar molecule
- weak forces
117
how do dipole-dipole forces occur
- polar molecules with permanent dipoles attract other polar molecules with permanent dipoles
- stronger than LDFs and dipole-induced dipole forces
118
how does hydrogen bonding occur
- when H is bonded to a hihgly electronegative atom (F, O, N etc.) to create a very polar bond
- this H atom forms strong electrostatic interaction (H bond) w electronegative atom of another molecule
119
relative strengths of intermolecular forces
LDFs < dipole-induced dipole forces < dipole-dipole forces < H bonding
120
volatility of covalent substances
covalent network: non-volatile cus of strong covalent bonds must be overcome
molecular covalent: volatile, intermolecular forces must be overcome, pretty easy
121
why do larger covalent molecules have lower volatility
- stronger LDFs leading to higher MP and BPs
122
electrical conductivity of covalent molecules
- both network and molecular covalent substances are not conductive -> lack delocalized e-s
- exception: graphite, silicon (semi)
123
solubility of covalent molecules
covalent network: insoluble cus strong covalent bonds
molecular covalent: soluble if intermolecular forces bw solute and solvent are stronger than bw solute molecules -> "like dissolves like" applies too
124
miscible vs immiscible
2 liquids that can form a solution when mixed in any proportion )ex. water + ethanol) vs liquids not forming a homogenous mixture
125
how do surfactants work
- have hydrocarbon chains with hydrophobic tail and hydrophilic head
- able to associate w both water and non-polar fats to wash grease away
126
chromatography
technique used to separate the components of a mixture
based on their relative attractions involving intermolecular forces to mobile and stationary phases
127
stationary and mobile phases of paper chromatography
stationary: chromatography paper
mobile: solvent
128
how does paper chromatography work
- paper places in chamber with small amount of solvent at bottom
- solvent moves up, components of mixture move diff distances based on relative affinities to solvent
- chromatography paper made of hydrated cellulose -> many -OH groups, very polar so attract water tightly
- components w greater affinities of solvent will dissolve more readily, travel further
129
thin layer chromatography stationary and mobile phases
stationary: rectangular plate made of glass or metal coated silica or alumina
mobile: solvent
130
how does thin layer chromatography work
- glass/metal rectangular plates are very polar from -OH groups
- polar substances travel slowly as they adsorb onto silica/alumina rather than dissolve in solvent and vv.
131
what are resonance structures/how do they occur
when there is more than one possible position for a double or triple bond in a molecule
- cannot be represented by a single lewis formula
132
how does delocalization occur -> in terms of resonance structures?
when e- are shared by 2+ atoms in a molecule/ion as opposed to being localized bw a pair of atoms
- can represent resonance structures using a single structure using concept of delocalization
133
where are e- most dense in benzene
- one e- for each carbon occupies p orbital which overlap above and below the benzene ring
- 2 donut like shapes of e- density above and below benzene
134
physical evidence for the resonance structure of benzene
- xray diffraction shows that all C-C are the same length -> length and strength is intermediate bw single and double bond
135
chemical evidence of the resonance structure of benzene
- doesn't like addition rxns as it would disrupt stabilizing effect of delocalized e- ring
- resonance increases stability of benzene molecule
- 1,2-disubstituted benzene compounds should be diff with alternating single double bonds, but are actually the same showing resonance
136
expanded octet
molecules containing atoms with more than eight valence e-
137
formal charge
charge an atom would have if all bonding e- in the molecule were shared evenly + if non-bonding e- were not shared at all
138
formula for calculating formal charge
formal charge=(number of valence e-) - [(number of nonbonding e-)+1/2(number of bonding e-)]
FC = VE - (NBE-1/2BE)
139
preferred conditions for formal charge to find most preferred Lewis formula
- as close to 0 as possible
- difference informal charges close to 0
- negative formal charges are assigned to the more electronegative atoms
140
why the strength of a double bond is not double the strength of a single bond
- single bond contains a sigma bond, double bond contains sigma and pi bond, a pi bond is generally weaker than a sigma bond
141
bond axis/internuclear axis
- invisible line bw the nuclei of 2 bonding atoms
- region of high electron density
142
sigma vs pi bonds
overlaps side by side for s orbital (circle) and p orbital (figure 8) vs. only for p orbital overlapping above and below bond axis
143
hybridization
mixing atomic orbitals to for new hybrid orbitals for bonding
144
why is hybridization energetically favourable (C as an example)
- promotion step absorbs energy BUT energy released from subsequent bond formation outweighs that
- promotion also doesn't require too much energy cus it relieves 2s electro on repulsion it experiences when paired
145
how are sp3 hybrid orbitals created (C as an example)
- one 2s e- is promoted to 2p
- the one 2s and three 2p atomic orbitals are combined producing 4 equivalent sp3 hybrid orbitals
- 25% s character and 75% p character
146
how are sp2 hybrid orbitals created (C as an example)
- combination of one 2s and two 2p atomic orbitals form 3 sp2 orbitals
- remaining unhybridized p orbital can form a pi bond w a parallel unhybridized p orbital on another atom
147
how are sp hybrid orbitals created (C as an example)
- combination of one 2s and one 2p orbitals forms 2 sp hybrid orbitals
- hybrid orbitals form sigma bonds, unhybridized orbitals can form pi bonds w neighboring atoms
148
hybridization of oxygen
electron configuration of 1s^2,2s^2,2p^4
- 2s and 2p orbitals hybridize to create 4 sp3 hybrid orbitals -> 2 lone pairs and 2 singly occupied hybrid orbitals
149
what is a metallic bond
the electrostatic attraction between a lattice of cations and
delocalized electrons
150
metallic bonding model
- when metallic atoms bond to other metallic atoms, their valence e-s become delocalized
- a "sea" of electrons that surround cations
151
conductivity of metallic structures
good electrical conductor
- can conduct electricity cus of delocalized e-s
good thermal conductor
- particles gain kinetic energy when heated, vibrations of cations increase-> vibrations easily passed to other cations as they're closely packed-> vibrations transfer energy to delocalized e-s, passing E to other parts of lattice
152
electrical resistance in metals
- as thermal energy increases, ions in lattice vibrate more, collide more to convert some kinetic energy converted into heat
- these increased collisions cause resistance
153
superconductors
materials that offer no resistance to electric current below a certain critical temp
154
malleability/ductility of metallic structures
good malleability
- when force is applied to metallic structure, layers of cations can slide past each other w/o breaking metallic bonds to surrounding deloc. e-
- cations can be rearranged easily
155
what doe the strenght of metallic bonds depend on
- ionic radius (larger ionic radius = decreased bond strength)
- ionic charge (greater ionic charge = increased bond strength)-> greater # of e- contributed to sea of e-
156
trend of melting/boiling points down group and across period (s and p block)
down group: ionic radius increases, metallic bond strength decreases -> lower MP/BP
across period: charge of cation increases, ionic radius decreases -> higher MP/BP
157
why are the MP/BP and conductivity of transition metals high?
- e-s in d sublevel and e- of outer level become delocalized -> greater e- density bw cations and e- sea
- larger # of delocalized e- also allow for good electrical conductivity
158
what makes substances brittle vs elastic
- brittle substances have atoms/ions that are unable to slide past each other
- metallic bonding is elastic, or rubber is elastic due to long polymer chains
159
plasticity
material retains deformed shape even after external force is removed
160
how does silicon exhibit both metallic and covalent character
- metalloid
- shiny like metal
- brittle, forms weak acidic oxide like covalent substances
- intermediate conductor
161
how does magnesium, oxide exhibit ionic and covalent character
- 50% ionic 50% covalent character
- unusually high lattice enthalpy
162
how does aluminium chloride exhibit both ionic and covalent character
- classified ionic but has significant covalent contribution
- unusually low melting point which doesn't match with place on bonding triangle
163
alloys
mixtures of a metal and other metals or non-metals
- they have enhanced properties
164
properties of alloys
- in pure metals, all cations same size, in alloys there are cations of different sizes, disrupting regular structure of lattice
- layers of cations cannot slide past each other as easily, alloys are stronger
165
polymers
large molecules, or macromolecules, made from repeating subunits called monomers (polymerization)
166
properties of polymers
- strong from macromolecular nature
- don't contain moving particles, low thermal and electrical conductivity
- high melting/boil points -> large # of intermolecular forces from length of polymer + need energy to untangle
167
crosslinking
when polymer chains are covalently bonded together -> forms dense covalent network
168
what characteristics make plastics useful
- has LDFs that are overcome thru heating, so can be reshaped easily
- good thermal insulator, used to make coolers
- strong covalent bonds so are chemically inert and durable
169
what makes recycling plastic difficult
- energy-intensive (plastic must be stored and cleaned)
- must be heated up but plastics release atmospheric pollution
- quality of plastics deteriorate when recycled
170
how does addition polymerization work
- breaking of double bond in each monomer
- when double bond in monomer breaks, allowing it to join with another monomer
171
how does condensation polymerization happen
reaction between functional groups in each
monomer with the release of a small molecule
172
carboxylic acid and alcohol as an example of condensation polymerization, what is formed
- forms polyester
- requires one monomer to have 2 carboxyl groups and other to have 2 hydroxyl groups
- ester link forms bw, water is released
173
carboxylic acid and amine as an example of condensation polymerization, what is formed
- forms polyamide
- requires one monomer to have 2 carboxyl groups and other monomer to have 2 amine groups
- amide link forms, water released
- can also be formed w one acid and one amino group to make amino acid
174
formation of proteins as a condensation polymerization
- are polypeptides
- monomers called 2-amino acids (have one side COOH and one side NH2)
- releases water when COOH and NH2 bond to make a peptide bond
175
hydrolysis
- reverse of condensation reactions where linkage formed by condensation is split using a molecule of water
176
how are biological macromolecules formed/break down
- all biological macromolecules form thru condensation rxns and break down by hydrolysis
177
periodicity of atomic radius
- from start to end of period, each atom has a slightly smaller atomic radius (1 more proton so higher nuclear charge)
- atomic radius increases greatly down a group as e- shielding increases
178
change of ionic radius when forming cations
- atom loses valence e- so more protons than e- so ionic radius is smaller than atomic radius (more e-s lost = smaller ionic radius)
179
change of ionic radius when forming anions
- atoms gains e- to fill valence sublevel som ore e- than protons
- force of attraction bw nucleus and e- are smaller, ionic radius is greater than atomic radius (more e-s gains = greater ionic radius)
180
isoelectric species
have the same e- configuration but formed from different elements w diff number of protons
- results in differently sized species
181
periodicity of ionization energy
- minimum energy needed to eject an e- out of a neutral atom
- decreases down a group as shielding effect is stronger, less energy needed to remove e-
- increases across a period as number of protons increases so outer e- are held closer by increased nuclear charge, more energy needed
182
periodicity of electron affinity
- energy released when an additional e- is attached to neutral atom/molecule
- more favourable the rxn=higher e- affinity
- F, Cl, Br have very high e- affinities as they're 1e- away forma full shell (inc. across perid)
- little known abt trend down a group
183
periodicity of electronegativity
- atom's ability to attract e- fron chemical bond to itself
- decreases down group: larger atoms will attract less e-
- increases across period: more valence e-= more attraction
184
trend of metallic character of group 1 alkali - reaction with water
- metallic character is closely linked to ionization energy (lower ionization energy=more metallic)
- alkali metals form bronsted-lowry bases when reacted w water
- ionization energy decreases down group 1 = metallic character increases = more vigorous rxn
185
trend of non-metallic character of group 17 halogens - reaction with halides
- non-metallic character closely linked to electronegativity-> decrease going down group = reactivity decreases
- the greater the EN diff bw two species, the more readily the reaction occurs
186
general equation of alkali metal oxides w water
M2O(s) + H2O(l) --> 2MOH(aq)
- many metal oxides are lewis bases, they donate an e- pair to hydrogen in water
187
general equation of group 2 oxides with water
MO(s) + H2O(l) --> M(OH)2 (aq)
- group 2 oxides are lewis bases, donate an e- pair to H in water
188
equations of non-metal oxides with water
- non-metallic oxides are lewis acids -> react with water to form other acids by accepting an e- pari from O in water
189
amphoteric species + example
- chemical species that behaves as both a lewis acid and base
- ex. Al2O3 (aluminium oxide)
190
periodicity of acidity/alkalinity of oxide elements
- across a period, oxides of elements are less basic and more acidic
191
rain is acidic/basic, why? how can you tell normal rain from acidic rain?
- acidic because as it falls, it dissolves CO2 from the atmosphere which is acidic
- carbonic acid is weak, can only reach pH of 5.6 so acid rain is when pH<5.6 <-acidity is cause from more than CO2
192
amphiprotic vs amphoteric
species that can accept and donate a proton (acting as acid and base) vs. species that can act as both acids and bases in general
193
oxidation state
a number assigned to an atom to show the number of electrons transferred in forming a bond
- the charge that atom would have if the compound were composed of ions
194
why is the oxidation state of an element always 0
- atoms of same elements have equal electronegativity so e-s are shared evenly across bonds
195
naming oxyanions
- oxyanions are polyatomic anions that include O atoms
- name polyatomic anion, then include oxidation state of non-oxygen atom as roman numeral in brackets
195
why are there discontinuities in ionization energy trends (ex. bw Be and B, N and O)
- Be + B: e- configuration in Be is 1s2,2s2 in B its 1s2,2s2,2p1 so even tho B has greater nuclear charge, the valence e- is at a higher sublevel which is shielded by 2s2 so less energy needed
- N + O: N has3 e- in 2p orbital who dont come into close proximity, paired e- in O creates increased repulsion making the e- easier to remove even tho it have greater nuclear charge
196
central atom with 5BP/0LP- electron group arrangement + shape
trigonal bipyramidal, trigonal planar and extra e added to top and bottom
197
central atom with 4BP/1LP- electron group arrangement + shape
seesaw, trigonal planar with one e removed and extra e added to top and bottom
198
central atom with 3BP/2LP- electron group arrangement + shape
T-shaped, 2 vertical e one horizontal e
199
central atom with 2BP/3LP- electron group arrangement + angles
linear, 180 degrees
200
central atom with 6BP/0LP- electron group arrangement + shape
octahedral, 4 e on horizontal plane and 2 e on vertical plane
201
central atom with 5BP/1LP- electron group arrangement + shape
square-based pyramid, 3e on horizontal plane and 2 e on vertical plane
202
central atom with 4BP/2LP- electron group arrangement + shape
square planar, 2e on horizontal plane and 2 e on vertical plane
203
central atom with 3BP/3LP- electron group arrangement + shape
T-shaped, 1 e on horizontal plane, 2 e on vertical plane
204
central atom with 2BP/4LP- electron group arrangement + angles
linear, 180 degrees
205
observations of lithium reacting with water
floats, reacts slowly, releases H, keeps shape
206
observations of Na in water
vigorous release of hydrogen, heat produced melts unreacted metal forming small ball that moves on water surface
207
observations of potassium in water
produces enough heat to ignite hydrogen release by reaction, produces lilac flame
208
properties of transition elements
- result of incompltee d-sublevel
- attraction bw delocalized e-s and metal ions are strong, high MP/BP
- most have unpaired d e-s -> creates magnetic moment which can align w external magnetic field -> transition metals are magnetic
209
why is zinc not a transition element
- doesn't form cations with an incomplete d-sublevel -> full d sub level so readily loses 2 e- from 4s orbital instead
- all d e-s in zine are paired in neutral state so not magnetic either
210
catalyst
species that reduces activation energy required for a reaction to occur, while not being used up
211
how are transition elements used as catalysts
- heterogeneous catalyst -> catalyst that exists in a separate phase to reactants
- reactants A and B adsorb onto active site of solid transition element
- reaction occurs on catalyst
- product desorbs
212
why do transition elements have variable oxidation states
- the successive ionization energies are close in value
- very small energy difference bw losing different #s of electrons
213
what are transition element complexes
- the coloured compounds transition elements form when bonded with ligands
214
ligands
molecules/ions with a lone pair of e- that can be donated to a transition element cation
- results in formation of coordination bond bw ligand and central cation
215
degenerate orbitals
orbitals that have the same energy
216
why are transition element complexes coloured
- orbitals split creating non-degenerate orbitals
- diff. in energy bw split orbitals allows e- in lower orbitals to be promoted as complex absorbs light
- colour observed is opposite to colour absorbed
217
what factors affect the colour of transition element complexes
- ligand and transition element -> same ligand can cause diff splitting depending on central metal cation -> stronger coord. bonds = greater energy gap bw split d orbitals
- diff charges on central cation will create diff splitting = diff colours
218
how to determine a ligand in a complex
- will be the molecules/ion directly bonded to the central atom, within the same square brackets
219
catenation
process where many identical atoms are joined tgt by covalent bonds, producing straight, branched and cyclic structures
220
functional group
atom/group of atoms that gives organic compounds their physical and chemical characteristics
221
saturated vs unsaturated hydrocarbons
all C-C bonds are single vs primary carbon chain must contained 1+ C-C double/triple bonds
222
aliphatic
to describe molecules that contain no aromatic rings
223
what are homologous series
- family of compounds, groups tgt based on similarities in structure and reactivity
- same general formula, differing by a CH2
224
general formula of alkanes vs alkenes vs alkynes
CnH2n+2 vs CnH2n vs CnH2n-2
225
general formula of halogenoalkanes
CnH2n+1X
226
general formula for alcohols
CnH2n+1OH
227
general formula for aldehydes and ketones
CnH2nO
228
generala formula for carboxylic acids
CnH2nO2
229
general formula for amines vs amides
CnH2n+3N vs -CONH2
230
primary vs secondary vs tertiary amines
N is bonded to 1 carbon vs 2 carbons vs 3 carbons
231
MP/BP trends in homologous series
MP/BP increases for each successive member of a homologous series cus longer chains = more contact points for intermolecular forces
232
structural isomers
compounds that have the same molecular formula but different connectivities of atoms
233
chain vs positional isomers
structural isomers w diff lengths of carbon chains vs structural isomers where position of functional group changes
234
primary vs secondary vs tertiary compounds
carbon atom (or N for amines) bonded to 1 other carbon atom vs C atom bonded to 2 other C atoms vs C atom bonded to 3 other C atoms
235
functional group isomers
structural isomers with atoms arranged differently to create diff functional groups
236
steroisomers
same molecular formula/connectivities/bonds but have diff spatial arrangements of atoms
237
conformational isomers vs configurational isomers
atoms rotating about a bond to form 2 possible orientations (possible when single bond present) vs can only change orientations by breaking and reforming bonds (typically double or triple bond present)
238
cis vs trans isomers
2 identical substituents are on the same side of carbon vs on opposite sides of the carbon
239
chiral carbon/stereocentre/asymmetric centre
carbon atom bonded to 4 different atoms/groups of atoms
240
optical isomerism
molecules with one or more chiral carbon atoms
- type of configurational isomerism
241
enantiomers
pair of optical isomers that are non-superimposable
242
racemic mixture (racemate)
50:50 mixture of 2 enantiomers
- does not rotate plane-polarized light
243
different chemical properties of enantiomers
- behave differently in chiral environments
chiral environments: non-racemic environment that distinguishes bw enantiomers
244
optical activity of enentiomers
- pair of enantiomers under same conditions will rotate plane polarized light by same angle in opposite directions
(+) enantiomer rotates light clockwise, (-) enantiomer rotates light counterclockwise
245
mass spectrometry
analytical technique to break up organic compounds into ions (fragmentation)
246
molecular ion/parent peak
peak with highest mass to charge ratio (m/z) (rightmost peak)
247
base peak
most abundant fragment
- shows up as relative abundance of 100
- most stable fragment formed -> easiest to break
248
what do the frequency of the vibrations of the bonds depend on in IR spectroscopy
- bond enthalpy: higher bond enthalpy = stronger bond = shorter wavelength = higher frequency
- mass of atoms: larger the mass = lower frequency of vibration
249
infrared spectroscopy
observes vibrations of bonds and evidence of the functional groups present
250
what makes a compound IR active
- molecule must experience a change in dipole moment during vibration (must have a dipole)
- hence diatomic molecules like Cl2, H2 etc. are IR inactive
251
proton nuclear magnetic resonance spectroscopy (H NMR), what does it tell us
determine number of chemical environments of a compound and number of H attached to an atoms
252
what does NMR need to work
only works with nuclei that exhibit nuclear spin (nucleus must have odd atomic # or odd mass #)
253
alpha spin vs beta spin
lower energy state vs higher energy state
254
hydrogen atoms in H NMR
- H nucleus behaves as a spinning charged particle like a small magnet
- 2 possible spin states are degenerate BUT when strong magnetic field is applied, the 2 states split into 2 diff energy levels (alpha and beta)
255
energy difference bw alpha and beta spin states
- proportional to the strength of the magnetic field
- as magnetic field increases, the energy difference increases
256
what is needed for a 1H nucleus to switch from alpha to beta spin
- for nucleus to be promoted, must be exposed to radio waves of energy equal to difference
- exact energy is diff for diff chem. environments -> shows as diff peaks on H NMR
257
what does it mean to be “in resonance”
when a nucleus absorbs RF radiation with energy equal to energy diff bw spin states (continual flipping)
258
what is the TMS
- maximum shielding “reference sample”
- Si bonded to 4 CH3
- silicon is less electronegative that carbon, methyl groups are relatively electron rich
259
characteristics that make TMS useful
- very stable and unreactive
- volatile liquid that can be removed easily and recovered
- readily available in pure form
- experiences maximum shielding
260
integration traces, what do they tell us
- area under a curve created by a peak on H NMR graph
- ratio tells us relative number of H in each chemical environment
261
why do peaks in H NMR split (spin-spin coupling)
- occurs when two diff types of protons are close enough that their magnetic fields influence each other
- may have magnetic fields that align or oppose external field (magnetic coupling)
262
upfield vs downfield in H NMR
frequencies closer to 0 vs. larger frequencies
263
number of protons causing splitting relationship to number of peaks
number of peaks will be one more that the number of equivalent protons causing splitting
264
when does magnetic coupling (spin-spin splitting) not occur
- between equivalent protons
- bw protons separated by 4 or more bonds
265
transfer of energy/matter in open vs closed vs isolated systems
transfer of matter and energy is possible across its boundary vs no transfer of matter, but transfer of energy may be possible vs. no transfer of matter or energy
266
temperature vs heat
state function so a change in value is independent of the pathway bw initial and final measurements (doesn't tell what happens bw initial and final) vs form of energy transferred from a warmer body to a cooler body
267
entropy and motion of particles at absolute zero
- all motion of particles theoretically stops
- entropy of system reaches minimum possible value
268
exothermic vs endothermic reactions
chemical rxn where heat is transferred from system to surroundings vs chem rxn where heat is absorbed into system from surroundings
269
sign of enthalpy change in exothermic vs endothermic rxns
negative enthalpy change vs, positive enthalpy change
270
relative stability of reactants and products in exothermic rxn
- reactants of rxn are at higher energy level, low stability
- products are at lower energy level, more energetically favourable
271
relative stability of reactants and products in endothermic rxn
- reactants at lower energy level
- products at a higher energy level, less stable
272
standard enthalpy change
heat transferred (change in temp) at constant pressure under standard conditions and states
273
specific heat capacity
amount of heat needed to raise the temp of 1g/kg of that substance by 1C
274
bond breaking is endo/exo while bond making is endo/exo rxn
endo, exo
275
bond enthalpy
energy required to break one mole of bonds (using homolytic fission) into one mole of gaseous covalent molecules under standard conditions
- results in formation of radicals
276
limitations of bond enthalpy
- average values so will differ from actual value
- don't account for intermolecular forces which are important in solids/liquids -> should only calculate in gas form
277
why are bond enthalpies only average values
- chemical environment changes upon removal of successive H atoms
- ex. if a molecule of methane had one H removed at a time, bond enthalpy for each would be different
278
hess's law
regardless of the route a chemical rxn proceeds, the enthalpy change will always be the same as long as initial and final states of system are the same
279
standard enthalpy change of combustion
enthalpy change that occurs when one mole of a substance in its standard state is burned completely in oxygen
280
standard enthalpy change of formation
energy change that occurs when one mole of a substance is formed from its constituent elements in standard states
281
values of enthalpy of formation for allotropes
- allotropes have different enthalpy of formation values
- most stable allotrope is the standard state and have enthalpy of formation of 0 (ex. carbon -> graphite will have Hf=0)
282
ionization energy
standard enthalpy change when one mole of e- are removed from atom/positively charged ions in gas state
283
enthalpy of atomization
standard enthalpy change that occurs when one mole of gas atoms of an elements is formed
284
