Terms in this set (311)
Outline the rxn between acids + reactive metals, metal oxides, metal hydroxides, hydrogen carbonates, and carbonates
1. Acid + Metal = Salt + Hydrogen gas (reactive metal - this is reason for corrosion)
2. Acid + Metal hydroxide = Salt + Water (neutralization - exo)
3. Ammonia + Acid = Salt (neutralization - exo)
4.Acid + Metal oxide = Salt + Water (neutralization - exo)
5.Acid + Carbonate = Salt + Water + Carbon
6. Acid + Hydrogen carbonate = Salt + Water +
Atoms of same element with same number of protons, different number of neutrons
Used to determine relative atomic mass
Vaporization -> Ionization -> Acceleration -> Deflection
Amount of deflection depends on mass and charge (smaller the mass, higher the charge -> greater deflection)
mass/charge ratio recorded, measures mass and the relative amount of all ions present
Use of radioactive isotope Carbon
Use of radioactive isotope of Cobalt
Use of radioactive isotopes of Iodine
I-131 and I-125
used as tracers in medicine for treating and diagnosing illnesses
orbitals with lowest energy are filled first
orbitals filled within the same subshell are filled singly first
First Ionization Energy
The energy required to remove one electron from an atom in its gaseous state. (kJ/mol)
X(g) -> X+(g) + e-
an element that possesses and incomplete d sub level in one or more of its oxidation states
have variable oxidation states, form complex ligands, have colored compounds, display catalytic and magnetic properties
increase the rate of a chemical reaction by providing a different reaction pathway and reducing the activation energy
Catalyst: Platinum and Palladium (Pt and Pd)
catalytic converters in cars
Catalyst: Iron (Fe)
Haber Process - Ammonia production
(3H2 + N2 <-> 2NH3)
Catalyst: Vanadium (V) Oxide
Contact Process - Sulfur Trioxide production
(2SO2 + O2 -> 2SO3)
(C2H4 + H2 -> C2H6)
Catalyst: Manganese (IV) Oxide
with Hydrogen Peroxide to form
Water and Oxygen
- Permanent magnetism
- Unpaired e- align parallel to each other in domains (even when no field present)
- spinning unpaired electrons create small magnetic fields
- line up with applied magnetic/electric field
- makes metal becomes weakly magnetic
- more unpaired e- the more paramagnetic
- all electrons in a transition metal complex are paired
A solution of known concentration
Ideal Gas Assumptions (Kinetic Theory)
- Gases are made up of molecules which are in constant random motion in straight lines
- There are no intermolecular forces between the gas molecules (or negligible)
- The volume occupied by molecules is much less than the relative volume of the container
- All collisions are perfectly elastic
PV = nRT Units
P - Pascals
V - cubic meters (divide decimeters by 1000 or centimeters by 1000000)
R- 8.31 J/Kmol
T - Kelvin
Why do real gases deviate from ideal gas behavior at low temperature and high pressure?
(VOLUME AND PRESSURE PROBLEM)
Assumption: Volume of molecules << volume of container (true at low pressure)
When pressure increases, proportion of total volume of molecules to the volume of the container INCREASES.
Small error (at low pressure) grows larger.
100 cm3 in 1000 cm3 -> 10%
100 cm3 in 250 -> 40%
Thus, volume of real gas at high pressure is larger than predicted by ideal gas law.
Why do real gases deviate from ideal gas behavior at low temperature and high pressure?
(PRESSURE AND FORCES PROBLEM)
Assumption: there are no intermolecular (attractive) forces between molecules. (wrong because condensation would not occur)
When gas at low pressure: particles widely spaced that IM forces cannot be 'felt'.
As pressure increases, the particles are pushed closer together - attractive forces have a greater 'felt' effect.
This reduces the pressure of the gas as the particles lose energy since they are attracted to other particles.
Low temperatures increase deviation - lower KE of particles increases strength of IM forces.
When does a gas behave MOST like an ideal gas?
Low pressure and high temperature
When does a gas behave LEAST like an ideal gas?
High pressure and low temperature
(these are the points at which the gas tends to change to a liquid)
region of space where there is a high probability of finding an electron
can hold two electrons of opposite spin
has a defined energy state for given e-config
Heisenberg's uncertainty principle
it is impossible to know exact position of an e- at a precise moment in time
Gamma > X-Ray > Ultraviolet > Visible > Infrared > Microwaves > Radiowaves
(decreases frequency, increasing wavelength)
c = f λ
c - speed of light (Constant)
f - frequency - the number of waves which pass a point in one second
λ - wavelength - distance between two successive crests
Energy of Photon
proportional to frequency of the radiation
E = hf
is equal to the energy change of an electron moving from different energy levels
When an electron transitions TO the first energy level, n = 1, from ANY other energy level.
Corresponds to the UV region of the spectrum (High energy)
When an electron transitions TO the second energy level, n = 2, from ANY OTHER energy level.
Corresponds to the visible region of the spectrum
When an electron transitions TO the third energy level, n=3, from ANY OTHER energy level.
Corresponds to infrared region of the spectrum.
Condensed Electron Configuration
Trends in first ionization energies as evidence for main energy levels (First 20 elements)
Ionization energy generally increases from left to right across a period, as the nuclear charge increases. As the electrons are removed from the same main energy level, there is an increase in the force of electrostatic attraction between the nucleus and outer electrons.
Ionization energy decreases down a group as a new energy level, which is further from the nucleus, is occupied. Less energy is required to remove outer electrons that are further from the attractive pull of the nucleus.
Trends in first ionization energies as evidence for sub energy levels (First 20 elements)
The Group 2 elements have the electron configuration ns2. The Group 13 elements have the electron configuration ns2np1. The electron removed when the Group 13 elements are ionized is a p electron. The electron removed when the Group 2 elements are ionized is an s electron. Electrons in p orbitals are of higher energy and further away from the nucleus than s electrons and so are easier to remove than electrons in an s orbital.
Group 15 elements have the configuration ns2npx1npy1npz1. Group 16 elements have the configuration ns2npx2npy1npz1. For Group 16 elements, the electron is removed from a doubly occupied 2p orbital. An electron in a doubly occupied orbital is repelled by its partner and so is easier to remove than an electron in a half-filled orbital.
How do the oxides change across a period?
Basic (Metals) -> Amphoteric -> Acidic (Non-metals)
Na - Al: Giant Ionic structures
P, S, Cl: molecular covalent
Si: Giant covalent
energy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous ions.
X (g) + e- → X
**property of ISOLATED gaseous atoms
ability of an atom to attract electrons in a covalent bond
**property of an atom in a molecule
Melting and Boiling points
must talk about INTERMOLECULAR FORCES
Alkali Metal with water
produce hydrogen and metal hydroxide
Li - floats reacts slow
Na - vigorous release of H2, heat melts metal, small ball moves around surface
K - more vigorously, heat ignites hydrogen produced - lilac colored flame
Displacement reactions of Halogens
F displaces Cl, Br, I
Cl displaces Br, I
Br displaces I
(Because Flourine is not as far away from nucleus - smaller radius)
Colors of Halogens
F - colorless
Cl - pale yellow
Br - orange/red
I - purple
Halide + Silver ions -> precipitate
Sulfuric (VI) Acid
SO3(l) + H2O(l) → H2SO4(aq)
Sulfuric (IV) Acid
SO2(g) + H2O(l) → H2SO3(aq)
Equation for chloric(VII) acid (HClO4)
Cl2O7(l) + H2O(l) → 2HClO4(aq)
Equation for chloric(I) acid (HClO)
Cl2O(l) + H2O(l) → 2HClO(aq)
Equation for formation of silicates
SiO2(s) + 2OH-(aq) → SiO32-(aq) + H2O(l)
Physical Properties of Transition Metals
- High melting point
- High electrical and thermal conductivity
- Malleable and Ductile
- High tensile strength
- Iron/Cobalt/Nickel are ferromagnetic
Chemical Properties of Transition Metals
- Variable oxidation states when forming compounds
- form variety of complex ions
- form colored compounds
- act as a catalyst
elements whose atoms have an incomplete d sub-shell or give rise to cations with an incomplete d sub-shell
When transition metal ions in solution have a high charge density and attract ligands which form coordinate bonds with the positive metal ion
a species that uses a lone pair of electrons to form a coordinate bond with a metal ion
uses lone pair of electron to form a covalent bond
number of coordinate bonds from the ligands to the central ion
have more than one available lone pair to form coordinate bond
complex containing at least one polydentane ligand
Why does the d-orbital split in a complex ion?
The electric field of a ligand (has lone pair of e-) causes repulsion of the electrons in the d orbital. As a result, the d orbital splits into two sub levels: three below and two above. The difference in energy depends on the ligand (water will cause less difference than cyanide) - this difference corresponds to specific wave length of visible light.
Why do complex ions exhibit different colors?
The energy difference between the split sub levels of the d orbital corresponds to a photon of visible light. As light shines on the solution, the electrons are excited and promoted and to the higher energy sub level, they absorb that discrete photon of light, exhibiting the complementary color that is not absorbed.
What does the color of the complex ion depend on?
- nuclear charge and identity of metal cation (greater nuclear charge = more ligand interaction = higher energy difference = higher energy photons absorbed)
- charge density of ligand (greater charge denisty = greater repulsion = greater splitting = higher energy difference = higher energy photons absorbed) (i.e ammonia stronger than water)
- geometry of complex ion (orientation and geometry of the ligand and d orbital - do they fit?)
- number of d electrons present/oxidation number of cation
electrostatic attraction between oppositely charged ions
(NOT!!!! the transfer of e-)
Physical properties of ionic compounds
MP/BP: Very high - electrostatic attraction between ions in lattice are strong
Volatility: low volatility
Solubility: in ionic or polar solvents but NOT in non polar solvents
Conductivity: when in liquid state (molten or aqueous solution)
They are brittle - movement puts same charge next to each other - repulsion causes split
the ease with which a solid/solute becomes dispersed through a liquid to form a solution
Difference in Electronegativity for ionic bond
1.8 units or more
Difference in electronegativity for covalent bond
much less than 1.8
Name all the polyatomic ions (including charge) (7)
electrostatic attraction between a shared pair of electrons and positively charged nuclei.
a group of atoms held together by COVALENT BONDS
(cannot use to refer to ionic bonded compounds)
Bond length and strength
with multiple bonds: length decreases and strength increases
Length - distance between two bonded nuclei
Strength - (enthalpy) energy required to break bond
- unsymmetrical/unequal distribution of electrons
- occurs when one atom is more electronegative than the other in a covalent bond
- behave like an ionic bond in some cases
Valence Shell Electron Pair Repulsion
Repulsion applies to domains - single, double, triple bonding pairs, or lone pairs
# of electron domains determines geo arrangement
shape determined by angles between bonded atoms
lone pair causes greatest repulsion - affect shape
different forms of an element in the same physical state
(different bonding gives rise to to different structures and properties)
- C atom is sp2 hybridized, covalently bonded to 3 others
- forms hexagons in parallel layers
- bond angle of 120
- layers held together by weak london dispersion forces (can slide over each other)
- good electrical conductor (one non bonded delocalized electron per atom)
- not a good thermal conductor unless heat is conducted parallel to crystal layers
- grey crystalline solid
- brittle, soft and slipper layers, high MP, most stable allotrope
- used as a dry lubricant, in pencils, and as electrodes in electrolysis
-C atome is sp3 hybridized and covalently bonded to 4 others.
- tetrahedral arranged in regular repeating pattern
- bond angle of 109.5
- non conductor of electricity (immobile electrons bc they are bonded)
- efficient thermal conductor (better than metals)
- transparent, lustrous crystal
- hardest known substance, brittle, high MP
- jewelry, tools and machinery for grinding and cutting glass
- C atom is sp2 hybridized, bonded to 3 others
- bonded to closed sphere of 60 carbon atoms
- 12 pentagons, 20 hexagons
- not a giant molecule (fixed formula)
- low thermal conductivity
- semiconductor at normal temp and pressure, easily accept e- to form anions
- yellow crystalline solid soluble in benzene
- light and strong, reacts with K to make superconducting crystalline material, low MP
- used as lubricants, nanotubes and nanobuds, uses as capacitors in electronics, and catalyst
- C atom covalently bonded to 3 others
- hexagon, bond angle of 120
- 2 dimensional material (Single layer)
- very good electrical and thermal conductor (one delocalized e- per atom)
- completely transparent
- thinnest material to exist (thickness of one atom)
- strongest material (100 x steel), flexible, high MP
- used in photovoltaic cell, touch screens, high performance electronic devices
SIlicone and silicone dioxide structure
- giant lattice structure as each atom is covalently bonded to four others in a tetrahedral arrangement
- SiO2 - empirical formula
- strong, insoluble, high MP, non-conductor of electricity
London dispersion forces
- temporary, instantaneous dipole formed by general movement of electrons
- induces a dipole in a neighboring atom
- weak forces of attraction occur between opposite ends of two temporary dipoles
- weak BUT increases with greater molecular size
- the only force between non-polar molecules
- in ALL molecules
-occurs between polar molecules that have a permanent separation of charge as a result of difference in electronegativity value
- The atom with larger EN value pulls electrons causing a permanent dipole
- positive and negative attract -> force
- strength depends on distance and orientation of dipole
- leads to solubility in polar solvents
- H bonded to N, O, F (high EN atoms)
- large EN difference results in N, O, or F pulling the electron pair away from hydrogen
- hydrogen, with no other electrons and no shielding, exerts a strong attractive force on a lone pair from the electronegative atom of a neighboring molecule
- strongest intermolecular attraction
Why are the melting and boiling points of covalent structures less than ionic structures?
- melting or boiling involves separating particles by overcoming the intermolecular forces between them (stronger the force = higher the MP/BP)
- overcome weak intermolecular forces that are easier to break than the electrostatic attractions in the ionic lattice
Melting point of macromolecular and giant covalent structures
- do not experience intermolecular forces
- covalent bonds must be broken for change of state to occur - NOT breaking of the forces
Solubility of large molecules
- polar bond is only small part of total structure
- solubility reduced as non polar part of molecule is insoluble
How do non-polar solutes dissolve in non-polar solvents?
- formation of london dipsersion forces between solute and solvent
How do polar solutes dissolve in polar solvents?
- dipole-dipole attraction and hydrogen bonding
Do giant molecular substances dissolve in solvents
- generally no
- too much energy needed to break strong covalent bonds in structure
Factors affecting strength of metallic bond
- number of delocalized electrons (greater = stronger)
- charge of cation
- radius of cation (smaller = stronger)
solid solution containing more than one metal, held together by metallic bonding
- has enhanced properties
- form because delocalized electrons are non directional
- lattice accommodates ions of different sizes
electrostatic attraction between a lattice of positive ions and delocalized e-.
Why are metals malleable and ductile?
movement of delocalized electrons is non-directional and random through a cation lattice, the metallic bond stays intact when while the conformation changes under applied pressure
Why are metals good electrical and thermal conductors?
- delocalized electrons are highly mobile and can move through metal structure when voltage is applied
- delocalized e- and closely packed ions enable efficient heat energy transfer.
Examples of Alloys (8)
Steel - Fe and C - high tensile strength but corrodes - used as structural material
Stainless Steel - Fe, Ni, Cr - high tensile strength, corrosion resistant - domestic and industrial appliances
Brass - Cu and Zn - Instruments
Bronze - Cu and Sn - coins, medals
Sterling Silver - Ag and Cu - Jewelry
- used to determine what the most stable structure is
- assumption: all atoms in a molecule or ion have the same electronegativity
FC = (# of valence e-) - (# of e- assigned to atom in lewis structure)
# of e- in lewis structure = (1/2 electrons in bonded pair + #of e- in lone pair)
reactive species with an unpaired electron
wavelength of light needed to break down oxygen
<242 nm (higher energy)
Wavelength of light needed to break down ozone
<330 nm (lower energy)
(200-325 nm) - corresponds to higher range of UV light UV-B and UV-C
Causes damage to living tissue
What damages the ozone layer?
- catalyze ozone destruction
- high energy UV breaks down the C-Cl bond (NOT C-F) to produce Cl free radicals
- Cl free radicals react with ozone and overall produce more oxygen and other free radicals
- NO2 breaks down in UV light to produce oxygen free rad
- oxygen radical reacts with ozone to produce oxygen molecule
- NO reacts with ozone to form NO2 and O2
- when two atomic orbitals on different atoms overlap along axis of nuclei.
- s + s
- s + p
- p + p
- when two p orbitals overlap above and below axis (side ways)
- two regions of electron density
- one 2s electron is promoted to 2p orbital
- 2s and 2p hybridize to form 4 new hybrid orbitals of the same energy
- tetrahedral arrangement - least repulsion position
- four equal sigma bonds are formed
- 2s electron is promoted to 2p orbital
- 2s orbital hybridized with TWO 2p orbitals to form 3 new hybrid orbitals of the same energy
- planar, 120 angle between them
- 3 sigma bonds and one pi bond (2 single bonds and a double bond)
- 2s orbital hybridizes with one 2p orbital to form two new linear sp hybrid orbitals (180 angle between them)
- the other two 2p orbitals overlap to form two pi bonds
- 2 sigma, 2 pi bonds
- one single bonds and a triple bond
atomic orbitals within an atom mix to produce hybrid orbitals of intermediate energy
Why can sulfur expand its octet whereas oxygen cannot?
S had a readily available empty 3d orbital which can be utilized, in O the d orbitals are too high in energy to be used.
ability to do work
the measure of the amount of heat energy contained in a substance.
- it is stored in the chem bonds and intermolecular forces as potential energy
-when substances react the difference in the enthalpy between reac and prod result in a heat change
heat is given out by the system to the surroundings
∆H is negative
bond in product is stronger than bond in reactant
heat is absorbed by the system from the surroundings
∆H is positive
bonds in reactant are stronger than bonds in product
standard conditions for enthalpy change
Pressure: 100 kPa
Concentration: 1 mol/dm3
Substances in their standard states
Standard state of substance
the normal, most pure stable state of a substance measured under standard conditions of 298 K and 100 kPa
specific heat capacity
the heat needed to increase the temperature by 1 K per unit mass of a material
standard enthalpy change of combustion
the enthalpy change for the complete combustion of one mole of a substance in its standard state in excess oxygen under standard conditions
- usually exothermic so ∆H is negative
Why are there discrepancies in the experimental and literature value for the combustion of ethanol? (assuming all heat from the reaction is absorbed by the water)
- not all the heat produced by the combustion reaction is transferred to the water (some needed to heat the calorimeter and some has passed to surroundings)
- combustion of ethanol is unlikely to be complete due to limited oxygen available
- experiment not performed under standard conditions.
the enthalpy change for any chemical reaction is independent of the route, provided the starting and final conditions, reactants and products are the same.
Standard enthalpy of formation
the enthalpy change that occurs when one mole of the substance is formed from its elements in their standard states under std conditions
(importance: measure stability of substance relative to elements, calculate enthalpy changes of all rxns)
specific heat capacity of water
Average bond enthalpy
the energy needed to break one mole of a bond in a gaseous molecule averaged over similar compounds
Density and specific heat capacity of aqueous solutions
Equal to water
Density - 1 kg/cubic meter
Shc - 4.18 kJ/kg/K
**this is a limitation, not true in real life
First Law of Thermodynamics
In chemical transformations, energy can neither be created nor destroyed
∆H = ?
∆H(prod) - ∆H(reac)
Limitation of bond enthalpy
- only valid for GASES (enthalpy for change of state would be needed)
- calculations are inaccurate because the intermolecular forces are not taken into account.
- environment of a bond not taken into account
Discuss bond strength in ozone relative to oxygen and its importance in the atmosphere.
- oxygen needs a higher energy (498 kJ/mol) since it has
double bond while ozone has 1.5 bonds (364 kJ/mol)
-formation and depletion of ozone - cycle
- starts with 1 O2 cleaved into two oxygen free radicals, by UV with λ<242nm.
- O2 (g) -UV→ O• (g) + O• (g)
- free radicals really want to react - lots of O2 to react w/
(exothermic) - makes stratosphere warm
- O2 (g) + O• (g) → O3 (g)
- What can this O3 do?
O3 (g) → O• (g) + O2 (g)
O3 (g) + O• (g) → 2O2 (g)
start with O2 again -Chapman cycle. This cycle keeps the concentration of ozone 10ppm
Lattice Enthalpy ΔHlat°
Lattice enthalpy is the energy required to disassociate/break down 1 mole of solid ionic compound into its gaseous ions. (Endothermic process)
NaCl (s) → Na+ (g) + Cl-
lattice formed: -ve H
broken: +ve H
magnitude depends on size of and charge of ions (smaller ionic radius and greater charge = higher the lattice enthalpy)
Enthalpy of Atomization
enthalpy change when one mole of gaseous atoms is formed from the element in its standard state under std
Na (s) → Na (g).
Enthalpy of hydration
enthalpy change when one mole of gaseous ions dissolves in sufficient water to give an infinitely dilute solution
(exothermic) - smaller ionic radius, greater charge = greater energy change
Enthalpy change of solution
enthalpy change when one mole of an ionic substance dissolves in water to give a solution of finite dilution
Na+ (g) + Cl-(g) → Na+ (aq) + Cl-(aq)
ΔH(sol) = ΔH(lat) + ΔH(hyd)
How does charge of the ions affect lattice enthalpy?
an ion with a higher charge will have a stronger electrostatic force of attraction therefore a higher lattice enthalpy
i.e: MgO vs NaCl
How does the ionic radius affect lattice enthalpy?
an increase in ionic radius results in a decrease of attraction between ions. Therefore, a larger radius = lower lattice enthalpy
i.e NaCl v KBr
What is the enthalpy change of formation for free elements?
Zero because they are not formed by anything. They are just elements.
i.e H2, O2
enthalpy change of combustion
enthalpy change when one mole of substance burns completely in oxygen under std.
- distribution of available energy among the particles in a system, the more ways the energy can be distributed the higher the entropy.
- measure of disorder of a system
- the more disordered = the more positive; ordered = more negative
if the overall transformation leads to an increase in total entropy (system + surroundings)
exothermic and entropy increases are spontaneous
Gibbs Free Energy (G)
relates the energy that can be obtained from a chemical reaction to the change in enthalpy, entropy, and temp
ΔGsystem = ΔHsystem - TΔSsystem
Absolute entropy of a substance
the entropy change per mole that results from heating the substance from 0 K to the standard 298 K
- absolute values can be measured unlike enthalpy
second law of thermodynamics
the direction of spontaneous change always increases the total entropy of the universe at the expense of energy available to do useful work
(total entropy of an isolated system can only increase over time)
order the states of matter in decreasing entropy
because the interaction forces prevent or
significantly decrease the distribution of energy
what does the entropy of surroundings depend on?
When ΔHsystem is exothermic, then the ΔSsurrounding should increase.
ΔHsystem is positive, ΔSsurrounding decreases since energy is taking from the surroundings.
ΔSsurrounding ∝ - ΔHsystem
How entropy is affected by temperature?
HIGH TEMP - small increase in entropy = negligible difference:
Let's say ΔSsurrounding is 100.
After whatever reaction, it has increased by 1
unit, hence 101.
101/100 = 1.01
Thus an increase in 1%.
LOW TEMP - small increase in entropy - massive difference:
Let's say ΔSsurrounding is 1.
After whatever reaction, it has increased by 1
unit, hence 2.
2/1 = 2.
Thus an increase in 100%!
ΔSsurrounding ∝ 1/T
Equation for entropy of surrounding
ΔS(surrounding) = (−ΔHsystem)/T
What does G tell us about spontaneity?
If G is -ve = spontaneous
If G is +ve = not spontaneous
(This is because -TΔS is the free energy that can do work, if it is greater than H, then there is more than enough energy and thus the reaction is spontaneous.)
What values of G result in a spontaneous/non spontaneous reaction?
not spontaneous if G is GREATER than + 30 kJ/mol
Spontaneous if G is LESS than -30 kj/mol
Equilibrium if LESS than 30kJ/mol and GREATER than -30 kJ/mol
Equilbrium and G
when G is at minimum and entropy is at maximum
Rate of Reaction
=(change of conc.)/(time)
Rate of Reactions Experiments
1. Change in volume of gas over time using a syringe or water displacement (when gas has low solubility in water)
2. Change in mass (not for light gases)
3. Change in light transmission - colorimetry - colored solutions in reactions:
Light source -> slit -> monochromator and prism -> one wavelength incident on sample -> detector (use control wavelength to compare)
4. Change in conductivity (solution with ions)
What does the rate of reaction depend on?
the probability that molecules will collide with sufficient energy AND proper orientation (so the reactive parts come in contact with each other)
The minimum energy colliding particles must have in order to have successful collisions leading to a reaction
- greater it is, the slower the rate as few molecules will have that energy
state between being reactants and products
Factors affecting rate
1. Number of collisions per unit time
2. increased chance of correct geometry in collision
3. Increased Kinetic Energy
Maxwell Boltzmann Distribution
- distributes the kinetic energy of the particles
- area under curve is constant
- curve shifts/broadens as more particles have a high velocity from an increase in k.e
How does a catalyst affect the rate?
by decreasing the Ea, it increases the rate without being chemically changed
- allows more collisions with the correct geometry to occur
What is the rate determining step?
The slowest step out of all reactions
- have a greater Ea
- catalysts are involved in RDS
the number of reactant particles taking part in that step.
- usually refers to RDS
What does the order of reaction indicate?
the number of particles taking place in the rate determining step with respect to a reactant
k, the rate constant
affected by temperature
units are determined from overall reaction order
Zero order reaction graph
for concentration v time
Zero order reaction graph for rate v concentration
First order reaction graph for concentration v time
First order reaction graph for rate v concentration
Second order reaction graph for concentration v time
Second order reaction graph for rate v concentration
What is A, the Arrhenius constant?
frequency factor /pre-exponential factor
- frequency of collisions and probability that they have the correct geometry
Is the rate constant, k, always constant?
No, only if the temperature is constant.
- as temp increases, more reactants have sufficient energy and therefore the rate increases
one species breaks down into two or more products
two species can collide with the necessary activation energy to give an activated complex
a chemical substance which cannot be isolated, consists of an association of the reacting particles in which bonds are in the process of being broken and formed - intermediate phase
- either breaks down to form products or reverts back to original reactants
When is equilibrium reached?
When the rate of forward reaction equals the rate of the backwards reaction
Chemical Equilibrium v. Physical Equilibrium
Chemical - decomposition of N2O4 in closed system
Physical - closed flask containing water - liquid water and water vapor
Characteristics of equilibrium
- occurs in closed system
- concentration remains constant (may not be equal)
- can be reached from either direction
- describes position of eqm.
- constant unless temperature changes
- eqm can lie closer to products or reactants
- less than 1 -> reac lies to left
- greater than 1 -> reac lies to right
Reaction Quotient, Q
measures the relative amount of products and reactants present during a reaction at a particular point in time
- eqm equation at NON EQM CONC.
- used to predict where reaction is going
- if Q = Kc -> eqm
- if Q > Kc -> Q will decrease and create more reactants
- if Q < Kc -> Q will increase and make more products
when there is a change in equilibrium, the reaction will change so that it experiences minimal chage
What happens to Kc if reaction is reversed?
What happens to Kc if the reactions coefficients are doubled?
square the value of Kc
What happens to Kc when reaction coefficients are halved?
Square root value of Kc
What happens to the Kc when two reactions are added together?
Multiply both Kc values
Kc1 x Kc2
Position of Equilibrium and Gibbs Free energy
If negative - spontaneous - produces more PRODUCTS
If positive - not spontaneous - more REACTANTS
If around 0 - equilibrium (entropy has max value)
Δ𝐺 = −𝑅T ln(𝐾)
- loss of one or more electrons from a substance
- loss of hydrogen atom (since one H atom has one e-)
- addition of oxygen to a substance
- the gain of one or more electrons
- gain of hydrogen atoms (one H = one e-)
- loss of oxygen
- oxidizes the other element and is therefore reduced
- reduces the other element and is therefore oxidized
Why do different oxidation numbers exist for transition metals and most main group non metals?
- transition metals have many oxi states due to electrons in d orbital
- non metals can change in oxidation number depending on the element it reacts with (electronegativity)
- shows which element is most reactive
- metal with high activity series = oxidizes easily (strong reducing agent)
- non metal with high activity series = reduces easily (Strong oxidizing agent)
- equivalence point indicator: color of the ions
1. Iron with Manganate (VII)
2. Iodine Thiosulphate
3. Winkler Method to calc dissolved o2.
'mind tool' - not true/arbitrary
- assumes that there is complete electron dominance in one atom - not shared
- +/- is before number
- the most electronegative atom will be reduced
elements that do not change during a reaction
How to identify what species is oxidized and what species is reduced
- write oxidation states and read how number has changed
Element's position on Activity Series and displacement
metals higher on the activity series can displace metal ions lower in the series from solution
Zn(s) + Cu2+(aq) -> Zn2+(aq) + Cu(s)
generates an electromotive force (potential difference) resulting in the movement of electrons from the anode(-) to the cathode(+) through an external circuit
- energy converted from spontaneous exothermic chemical processes to electrical energy
Characteristics of Voltaic cells
- 2 metal electrodes submersed in two separate solutions containing ions of the metal electrode
- metal with highest activity series oxidizes (at anode) -> e- travels to other metal
- metal ions in solution at cathode are reduced as they gain the e-.
- current only produced when there is a difference in reactivity
What is the role of the salt bridge?
- completes circuit
- keeps half cells electrically neutral as anions and cations accumulate in their respective cells
Metal | Oxidized metal || Metal ion | Reduced metal
i.e: Zn(s) | Zn2+(aq) || Ag+
(aq) | Ag(s)
|| - salt bridge
no need to balance
Which direction do the electrons flow?
From anode to cathode
Standard Electrode Potential, E
the potential difference of the reduction half-equation under standard conditions measure relative to the SHE.
std conditions: solute concentration - 1 moldm-3 or 100kPa (for gases)
Standard Hydrogen Electrode (SHE)
- H2 gas at 1 atm
- 298 K
- H+ solution at 1M
- Platinum electrode (inert)
Gives standard potential only if same temperature, pressure, and concentration of solution
- Std E potential: +ve or -ve -> depends on the direction in which e- are flowing (anode to cathode - which is being oxidized/reduced)
-ve: if element is good at oxidizing -> element to hydrogen
+ve: if element is good at reducting -> hydrogen to element
make non-spontaneous redox rxns occur by providing energy in the form of electricity from external source
- converts electrical energy to chemical energy
Characteristics of Electrolytic cells
- power source (long line positive, short line negative)
- INERT positive anode and negative cathode
- electrolyte (molten or aqueous)
-oxi at anode(+) and red at cathode (-)
Differentiate between electron and ion flow in a voltaic cell
- electrons flow through electrodes and wire
- ions flow through salt bridge
Differentiate between electron and ion flow in a electrolytic cell
- electrons flow through wire electrode (everything)
- ions flow to respective electrodes
the highest positive E (potential) will be most spontaneous and will therefore react
(Because G = -nFE -> most positive E will lead to -G which is spontaneous)
How does position on electrochemical series affect ion discharge?
the more positive the E value is, the more spontaneous, and therefore will be discharged at the electrode
How does the concentration affect ion discharge?
If there is a greater concentration of ions it will be discharged at the electrode when the electrode potential of water ions and the aq solution ions are close
How does the nature of the electrode affect ion discharge?
(when electrodes are not inert)
some electrodes oxidize/reduce in place of ions.
i.e Copper electrode in sol of copper sulfate - positive electrode is oxidized and releases electrons to form copper (II) ions - deposited at negative electrode conc. of solution stays constant throughout electrolysis
Factors that affect the amount of product formed in electrolysis
- duration of electrolysis (s)
- current (A) (if doubled - twice as many electrons flow - twice as much product)
- charge on ion (one mole of X+ requires one mole of e- to be formed where one mole of Y2+ requires two moles of e- - same current passes through so more of X+ is made)
**1 mole of e- has charge of 96500C
How to calculate the amount of product formed in electrolysis.
1. Find the amount of charge - Q = It
2. Find amount of charge needed in chemical equations (i.e NaCl: Na+ + e- -> Na; Cl- -> Cl +e-) moles of Na/Cl = (Total charge of circuit found in 1)/(total charge of Na/Cl found by mol of e- by 96500)
3. mass = mol
Mr OR volume = mol
molar vol of gas
Calculating electrode potential
- the more positive value will be spontaneous (G = -nFE)
- flip the more negative half equation
- add both half equations
- electrode potential is not per mole - do not multiply it
How do you a reaction is spontaneous by using E values
- more positive is spontaneous
- negative is not spontaneous
G = -nFE
(if 0 - equilibrium - no potential different - no work done)
Physical changes during electolysis
1. Color change - trans metal in sol is changing
2. precipitations - when there is an accumulation of metal at cathode
3. Effervescence - when non metals are oxidized into elements
4. pH change - OH- or H+ formation
the electrolytic coating of an object with a metallic thing layer
- cathode - the object we want to plate
- anode - the metal we will plate with
- electrolyte - contains ions of metal we will plate with
anode dissolves and oxidizes to become ion -> the ions are reduced at the cathode and is deposited on the electrode
- this process can be used to purify metals if impure metal is used as anode.
cells in series
- same amount of electricity passes through cells
- look at each cell individually - find molar ratios by looking at moles of electrons
must have lone pair to accept the H+
Conjugate acid-base pair
pair of species differing by one proton
c. acid - species formed after the base has accepted a proton
c.base - species formed after acid has lost proton
species that can act as both bronsted-lowry acids and bases
- must have H+ that can be released and lone pair to accept H+
e.g: water, hydrogencarbonate
(amphoteric - broad term for ALL molecules that act as acids and bases, amphiprotic applies only to loss and gain of H+)
How is a proton aq solution represented?
H+ or H3O+
How should the location of the proton transfer be indicated?
- compounds contain more than one H - location of ACIDIC H atom must be shown
- ethanoic acid: CH3COOH instead of C2H4O2 so that conjugate base can be shown
as CH3COO- instead of C2H3O2- to show position
- can tell whether sol is acidic or basic
- made of weak acids
Acid - Red
Alkali - Blue
Methyl Orange Indicator
Acid - red
Alkali - Yellow
Acid - Colorless
Akali - pink
- determines exact conc of acid or alkali by adding increments of alkali or acid (respectively) up until end point (where solution is neutralized)
Effect of a change of one unit in pH
10-fold change of hydrogen ion concentration - logarithmic scale
-pH 1 is 10x more acidic than pH 2
pOH - used for concentration of OH- ions
- high conc of OH- = very basic
- very basic = low pOH
- inverse to pH
ARTIFICIAL scale used to distinguish between acid, neutral, and basic/alkaline solutions
What is the ionic product constant of water
𝐾w = [H+][OH-] = 10^-14 at 298K
[H+]=[OH-] so take sqrt of 10^-14
- [H+]=[OH-]=10^-7 ; thus pH and pOh is 7 - neutral
digital indicator that does not rely on color - finds out quantitative value by sensing conc of H+ through electrodes
combination of a variety of indicators so that a color corresponds to a specific pH number
- completely dissociates into ions
- they become fully ionized
- this is a one way reaction
HCl -> H+ + Cl-
NaOH -> Na+ + OH-
- do not completely dissociate into ions
- partially ionized
- exists in equilibrium
CH3COOH -> CH3COO- + H+
NH3 -> NH4+ + OH-
Why do strong acids/bases have a higher conductivity than weak acids/bases when the concentrations are equal?
- strong/weak -> defines how they dissociate
- ions conduct electricity
- stronger acids/bases have more ions than weaker acids/bases and therefore conduct more electricity
"strong/weak" vs concentration
- acid strength is an inherent property of the molecule
- conc is how much mol per vol
-weak acid with high conc and strong acid with low conc can have the same pH.
- good proton donor
- has a weak conjugate base because the rxn will not reverse to produce the acid bc it is bad at accepting H+
(weak acids have a strong conjugate base that pushes eqm to left)
- good proton acceptor
- has a weak conjugate acid bc it does not reverse to produce base since the c.acid is bad at donating H+
- rxn is in one direction only
Strong Acid examples
1. Hydrochloric acid (HCl)
2. Nitric acid (HNO3)
3. Sulfuric acid (H2SO4)
4. Hyperchloric acid (HClO4)
Strong Base examples
All group 1 hydroxide. Ex,
4. Barium hydroxide, Ba(OH)2
How to distinguish between a strong and weak acid/base?
1. Conductivity - strong acid/base will be greater because more ions - measured by pH or conductivity meters.
2. Rates - strong acid/base will have faster rate of rxn w/ metals, metal oxides etc.
3. pH - indicate how much H+ is present - only works when solutions are equimolar
monoprotic v diprotic
monoprotic acid produces one mole of hydrogen where a diprotic acid produces two moles of hydrogen
(remember for pH calculations the H+ conc is two times H+ for diprotic acids)
Is rain naturally acidic or basic?
Acidic because of the dissolved CO2 content - produces weak carbonic acid. It has a pH of 5.6.
This is natural because CO2 is released by living organisms
pH below 5.6
process by which acidic particles/gases precipitate/leave the atmosphere
Wet - acidic oxides reacting with water in air - rain, snow, fog etc.
Dry - acidic gases, particles
How is acid deposition formed? (Sulfur)
1. Sulfur oxides
- sulfur dioxide - naturally from volcanoes and industrially from combustion of fossil fuels and extraction of sulfide ores
(S + O2 -> SO2)
- in sunlight sulfur dioxide is oxidized to sulfur trioxide
(SO2 + 1/2O2 -> SO3)
- Oxides react with water in the air to form sulfurous acid and sulfuric acid
(SO2 + H2O-> H2O3) and (SO3 +H2O -> H2SO4)
How is acid deposition formed? (Nitrogen)
2. Nitrogen oxides
- naturally: electrical storms and bacterial action
- industrially: NO is produced in the internal combustion engines.
(N2 + O2 -> 2NO)
- this can be oxidized to nitrogen dioxide in the air
2NO + O2 -> 2NO2
- Nitrogen dioxide reacts with water to form nitric acid and nitrous acid
(2NO2 + H2O -> HNO3 + HNO2)
- or Nitrogen dioxide oxidized directly to nitric acid by oxygen in the presence of water
(4NO2 + O2 + 2H2O -> 4HNO3)
Impact of acid deposition
- Leaching - metals can react with the acid - become soluble and wash away - i.e Mg2+
- no Mg2+ can cause reduction in chlorophyll - stems photosynthesis
- Aluminum in rocks react with acid and mix with soil water - Al3+ ion damages roots and prevent tree from taking up water and nutrients
2. Lakes and Rivers
- nitrates lead to eutrophication
- high levels of Al ions can kill fish
- aquatic life sensitive to pH (especially below 6, dead if below 4)
- erodes stone (marble and limestone is eroded)
4. Human Health
- inhalation of water vapor containing acids increases risk of asthma, bronchitis also deteriorated trachea and bronchus
- High levels of metals in water
Pre-combustion methods to reduce acid rain
- removing sulfur before burning it by sinking heavy sulfides or making the sulfur react with hydrogen to then capture the gas
post-combustion methods to reduce acid rain
- removing sulfur after it burns by precipitating it to calcium sulfate
- liming of lakes
Lewis Acid and Lewis Base
Acid - electron pair acceptor
Base - electron pair donor
**always forms a coordinate bond
(if not PAIR - redox rxn)
Nucleophile and electrophile
nucleophile - excess of e- -> base
electrophile - lack of e- -> acid
Explain why the Kw is temperature dependent and state the effect of the pH
Kw is an equilibrium constant
- temperature has an effect
- dissociation of water is bond break -> endothermic -> when temp increases it favors the dissociation - have more H+ and thus a lower pH
- temperature must always be stated with pH
-pH changes but ACIDITY does not as ratio of dissociation of H+ and OH- is the same
Ka - dissociation of WEAK acid
- weak acid - equilbrium so Ka is the equilibrium constant
HA(aq) + H2O(l) ⇌ A-(aq) + H+(aq)
Ka * [H2O] =[𝐴-][𝐻+]/[𝐻𝐴]
Assuming H2O is constant.
- endothermic so greater temperature means more dissociation and therefore a stronger acid
(in calculations: if acids are very weak the eqm conc of acid can be assumed to be same as initial conc)
Kb - dissociation of WEAK base
- weak base - equilibrium so Kb is the eqm constant
B(aq) + H2O(l) ⇌ BH+(aq) + OH-(aq)
Ka * [H2O] =[BH+][OH-]/[B]
Assuming H2O is constant.
-endothermic so greater temperature means more dissociation and therefore a stronger base
For a conjugate acid base pair, Kw = ?
Ka x Kb = Kw
reinforces conjugate acid base pair - when one is strong the other needs to be weak
pKa + pKb = 14
value of Ka/Kb for weak acid/base
- the weaker the weak acid/base the smaller the value for Ka/Kb BUT the larger the value of pKa/pKb
pH and pOH represent values that are...
given at equilibrium
(thus H+ and OH- concentrations are at eqm)
- a weak acid/base that is resistant to pH changes when SMALL amounts of acids and bases are affed
- do not have unlimited capacity
Strong acid and Strong Base pH Curve
e.g. HCl + NaOH
- initial pH is low because strong acid (i.e HCl)
- sharp increase from pH 3 to 11 -> equivalence point midway
- curve flattens at high pH because strong base (i.e NaOH)
- equivalence point: pH 7
Weak acid and Strong Base pH Curve
e.g CH3COOH + NaOH
-Initial pH is low because weak acid (i.e CH3COOH)
- pH stays relatively constant at beginning - buffer region because conjugate weak acid and strong base are the same concentration at half equivalence point
- sharp increasefrom 7 to 11
- the curve flattens at high pH because it is a strong base (i,e NaOH)
- the equivalence point is greater than pH 7
Strong acid and weak base pH curve
e.g HCl + NH3
- initial pH is quite high because weak base (NH3)
- pH is relatively constant at beginning - buffer region because concentration of conjugate strong acid and weak base is the same at half the equivalence point
- curve flattens at low pH because strong acid (HCl)
-equivalence point is less than pH 7
when stochiometrically equivalent amounts of acids and bases have reacted so the solution only contains salt and water.
- pH of equi point depends if salt is acidic or basic (See salt hydrolysis) i.e weak acid and strong bases has equi point greater than pH 7
represents the region where small additions of acid or bases result in little or no change in pH
if you look at slope of pH curve you can see that the buffer region's gradient is not steep but gradual
Half equivalence point
point where half the acid has been neutralized by base and converted to salt
- here the pH i= pKa or pOH = pKb
- to find: divide volume at equivalence point by 2 to find 1/2 equi point volume and then use graph to find pH
reaction of cation or anion with water which ionizes water molecule into H+ or OH-
- reacting a weak acid and strong base so that there are equal concentrations of the weak acid and its salt at eqm
e.g CH2COOH +NaOH -> NaCH3COO + H2O
CH2COOH =1 mol; NaOH = 0.5 mol; NaCH3COO = 0; H2O = 0
(((CH2COOH = 0.5 mol))), NaOH = 0 mol, (((NaCH3COO = 0.5 mo)))l, H2O = 0.5 mol
- solution of weak acid and its salt
e.g NaCH3COO -> Na+ + CH3COO-
CH3COOH <-> CH3COO- + H+ (eqm lie
s to left)
- high concentration of CH3COO- and CH3COOH
- acid added H+ + CH3COO- <-> CH3COOH
- base added: OH- + CH3COOH <-> CH3COO- + H2O
series of compounds of the same family, with the same general formula, which differ from eachother by a common structural unit
(chemical properties are the same, gradual change in physical properties)
compounds with the same molecular formula but different arrangements of atoms
reactive part of molecule
class: alcohol functional group: hydroxyl -OH
contain only single bonds
react by substitution
contain double or triple bonds
react by addition
resonance hybrid with delocalized e-
Kekule's Benzene Model
three double bonds
How do we know that Benzene is a resonance hybrid with delocalized e->
1. The C-C bond lenghts are all the same (0.140 nm) -> lies between value for C double (0.134nm) and single bonds (0.154 nm)
2. The enthalpy of hydrogenation of cyclohexene is -120 kj/mol. Thus, the expected enthalpy of H for benzene would be 3 times that (-360kj/mol) -> Experimental value is (-210 kj/mol)
3. There is only one isomer of benzene
4. Undergoes substitution reactions more often (expected to be addition because the double bonds would break)
trend of boiling points in homologous series
^ carbons in chain = ^ london forces = ^ more energy needed to break down force = ^ greater boiling point
The effect of molecular shape on the boiling point (in homologous series) i.e branching
- Branching occurs in secondary and tertiary compounds - molecules are more spherical in shape - this reduces the contact surface area between them - reduces bp
Trend in solubility of alcohols in water
- the larger the hydrocarbon chain - the less soluble because only the -OH group is polar where the chain is not
Combustion or Substitution (with halogens in UV light)
- because strong C-C and C-H bonds, unsaturated, low polarity/solubility
- low reactivity thus undergo free radical substitution
Combustion of Alkanes
Reaction is very exothermic
(C=O and O-H is much stronger than C-C and C-H)
incomplete combustion occurs when lack of oxygen to produce with CO or soot
Substitution rxn - alkanes
- UV light required
- forms halogenoalkane
free radical mechanism - initiation
- Cl2 splits heterolytically into Cl• + Cl• (two free radicals) - highly reactive
- requires UV light to occur
free radical mechanism - propagation
- Chlorine free radical, Cl•, reacts with other compound that does not contain free radical i.e methane CH4
- CH4 + Cl• -> CH3• + HCl
- produces another free radical - continues the mechanism
- CH3• + Cl2 -> CH3Cl + Cl•
- this triggers a chain reaction and continues for a long time
- the free radical is on the C/Cl not the H
free radical mechanism - termination
- two free radicals react with eachother to form a stable compound
- Cl• + Cl• -> Cl2
Alkene to Alkane
- Ni Catalyst
e.g. CH2CH2 + H2 -(Ni, 150C)->CH3CH3
Alkene to dihalogenoalkane
- addition of diatomic halogen
- Bromination - yellow/orange to colorless
CH2CH2 + Br2 -> CH2BrCH2Br
Alkene to halogenoalkane
CH2CH2 + HCl -> CH3CH2Cl
(Reactivity: HI > HBr > HCl because HI has weakest bond due to shielding in I)
Alkene to Alcohol
- heat with steam
- concentrated sulfuric acid (H2SO4)
- step with intermediate
-> CH3CH2HSO4 -(H2O)-> CH3CH2OH + H2SO4
Alkene to Poly(alkene)
- addition polymers made up of individual monomers
How can you test for the presence of an alkene group?
- bromine will go from yellow/orange to colorless
- if there is plenty of oxygen: forms carbon dioxide and water
- if lack of oxygen: forms either CO or C (soot) and water
Oxidation of Alcohol (generally)
- rxn with oxidizing agent: Cr2O7-
- Cr(VI) is reduced to Cr(III)
- Color changes from yellow to green when reduced
- the alcohol with this mixture is oxidized when heated
- rxn pathway depends on classification of alcohol
Oxidation of primary alcohol
- two step reaction
Alcohol -(1)-> Aldehyde-(2)-> Carboxylic Acid
+ Cr2O7- (oxidizing agent)
(after aldehyde formed, it can be removed through distillation since it has a lower BP than Alc and Carboxy Acid)
leave in same conditions as above
however heat under reflux to form carboxylic acid
e.g: CH3CH3 ---(Cr2O7-)---> CH3CHO ---(Cr2O7)---> CH3COOH
Cr2O7- changes from yellow to green when reduced
Oxidation of secondary alcohol
- one step reaction
Alcohol -> Ketone
+ Cr2O7- (oxidizing agent)
+ heat under reflux
- forms ketone and water
e.g Propan-2-ol ----(Cr2O7- and reflux)----> propanone + water
Cr2O7- changes from yellow to green when reduced
oxidation of tertiary alcohols
does not undergo oxidation under similar conditions, however if it does undergo oxidation its carbon skeleton is broken down
- when Cr2O7- is added, there is no color change
- organic reactions can be VERY slow
- heating is often used to increase the rate of reaction
- many organic compounds have low boiling points and will vaporize upon exposure to high heat
- preventing the reaction from proceeding in full.
- heating under reflux is used - heating a solution with an attached condenser to prevent reagents from escaping.
Alcohol + Carboxylic Acid <-> Ester + H2O
Water is produced - so 'condensation'
What is needed: conc H2SO4 (catalyst) and carboxylic acid
The equation is at equilibrium
Ester has the lowest boiling point ( no H bonding) thus it can be separated by distillation
- Ester is very non polar and is insoluble in water so forms a precipitate - layer on surface of water
e.g: methanoic acid + ethanol -> ethyl methanoate
Is an ester soluble?
no, it is largely non polar and forms a precipitate on the surface of water
Characteristics of halogenoalkanes
- they are saturated - therefore, undergo nucleophilic substitution
- for primary halogenoalkane: SN2 mechanism
- for tertiary halogenoalkane: SN1 mechanism
- they have polar bonds (halogens are highly electronegative) -> more reactive
- "Nucleus seeking" - it is electron rich - usually has lone pair that it can donate
- nucleophile replaces the halogen to form halogen anion
- i.e OH- replaces Cl -
Nucleophilic substitution of primary halogenoalkane
- SN2 Mechanism - bimolecular neucleophilc substitution
- one step with a transition state (carbocation intermediate)
- stereospecific because bond breakage occurs before bond breakage - an inversion occurs (think umbrella on windy day)
- Rate is dependent on both the Nucleophile and halogenoalkane since the RDS is before the intermediate - this rate = k[OH-][CH3Cl]
- the carbon halogen bond breaks heterolytically, giving Cl both electrons
Why is the SN2 mechanism described as stereospecific?
- 3D arrangement of reactants determines 3D configuration of products
- bond formation occurs before bond breakage (carbocation intermediate bonded to both the halogen and nucleophile before halogen becomes ion)
What type of solvents do SN2 reactions favor?
POLAR APROTIC SOLVENTS
- they cannot form hydrogen bonds because they do not contain -OH or -NH bonds
- they may have strong dipoles
- they solvate the metal cation (Na+) instead of nucleophile (OH-)
- Therefore the OH- is an unsolvated bare nucleophile and has a higher energy state - thus it increases the rate of reaction
- e.g. propanone and ethanantrile
they cannot form hydrogen bonds because they do not contain -OH or -NH bonds
they may have strong dipoles
they solvate the metal cation (Na+) instead of nucleophile (OH-)
Therefore the OH- is an unsolvated bare nucleophile and has a higher energy state - thus it increases the rate of reaction
Nucleophilic substitution of tertiary halogenoalkane
- SN1 Mechanism - unimolecular neucleophilic substitution
- 2 step reaction pathway - first step is the RDS - only the concentration of the halogenoalkane is taken into consideration -> rate = k[(CH3)3Cl]
- only happens in tertiary halogenoalkane because steric hindrance makes it difficult for a nucleophile to attack carbon
1. halogenoalkane ionizes because the other alkyls 'push it out'
2. The C-Cl bond breaks heterolytically, giving both e- to Cl to form Cl-
3. Halide ion is lost and thus the carbon has a partial and temporary +ve chage forming the carbocation intermediate
4. the carbocation intermediate is stable because the 3 alkyl groups around C create the positive inductive effect - electron donating effect
5. this allows space for the nucleophile to attack the partially positive carbon - since the carbocation has a planar shape, the nucleophile can attack from any direction - thus it is not sterospecific
What is the positive inductive effect?
- the alkyl groups have an electron donating effect
- this balances the partial positive charge of the intermediate carbocation
- thus the carbocation is stabilized until nucleophile comes
What solvents do SN1 reactions favor?
POLAR PROTIC SOLVENTS
- they contain either -OH or -NH bonds - form hydrogen bonding
why are they favored?
- the carbocation needs to be stabilized
- the protic solvent is effective in stabilizing it by solvation involving ion-dipole interaction
good protic solvents?
- Carboxylic Acids
Benzene reaction pathways
- delocalized e- allows for a stable compound
- addition reactions would break the double bonds and thus reduce stability - the products would end up with a greater energy
- Benzene undergoes electrophilic substitution
- electrophile is electron deficient and is attracted to the electron cloud of benzene
Benzene + Cl2 -> Chlorobenzene
Benzene + HNO3 -> Nitrobenzene
electron pair split evenly into two radicals
electron pair taken by one atom to form an ion
What is the difference between in amine and amide?
Amine - NH2 functional group
Amide - Carboxyl AND the NH2 functional group (on same carbon)
What is the difference between carbonyl and carboxyl?
Carbonyl (of the ketone): -CO
Carboxyl (carboxylic acid): -COOH