IB Chemistry

Vapour pressure
Pressure exerted on a liquid when rate of evaporation equals rate of condensation
Atmospheric pressure
Pressure exerted on a liquid
Boiling point
Vapour pressure equals atmospheric pressure
Lewis theory of acid
Acid = electron acceptor
Base = electron donor
Bronsted Lowery
Acid = proton donor

Base = proton acceptor

Arrhenius
Acids produce hydrogen ions
Bases produce hydroxide ions
Buffer
Resistant to changes in pH on the addition of small amounts of acid or alkali
Standard electrode potential
Emf generated when the half cell is connected to the standard hydrogen electrode by an external circuit and a salt bridge, under standard conditions.
Standard hydrogen electrode
298K
Hydrogen gas at 1 atm
Platinum electrode
Solution of 1M H
E=0.00
Negative electrode potential
Bad at gaining electrons
Oxidation
Always at Anode
Losing electrons
Reduction
Always at Cathode
Gaining
Positive electrode potential
Greater tendency to be reduced / gain electrons than hydrogen.
Operations of a mass spectrometer
Vaporisation
Ionisation: positively ionised so it can be accelerated by electric field. And so it can be deflected.
Acceleration: by oppositely charged plates so they all have the same kinetic energy
Deflection: deflected by magnetic field according to mass: charge ratio
Detection: electronically
Conditions in a mass spectrometer?
Low pressure / vacuum
The more reactive a metal, the more … it’s electrode potential
Negative
Salt bridge
Contains an aqueous solution of ions that are enables negative charge to be carried in the opposite direction to that of the electrons (from anode to cathode).
Standard conditions for measuring electrode potentials (5 specs)
Solutions have a concentration of 1 mol dm-3
All gases have a pressure of 1 atm
All substances used must be pure
Temperature is 298K
If the sol half cell does not use a solid metal then platinum is used as the electrode
Ester
COOR
Negative electrode potential
Less tendency to be reduced than the hydronium ion.
Cell potential
Difference in the tendencies of the half cells to be reduced.
E(where reduction occurred) – E(where oxidation took place).
Gibbs free energy is equal to INCLUDE BOTH EQUATION
= T – HS
= – Cell potential
Spontaneous reaction G is…
Negative
Structural isomers
Same molecular formula but with different arrangements of atoms in space.
Alcohol
OH
Aldehyde
COH
Ketone
RRCO
Carboxylic acid
COOH
Amine
NH2
Amide
CONH2
Nitrile
CN
Stereoisomers
Same molecular formula and structural formula but differ in the three dimensional orientation of the atoms in space
Primary alcohol/ halogenolkane
Carbon it is attached to is attached to one other carbon
Secondary alcohol/ halogenolkane
Carbon it is attached to is attached to two carbon atoms
Tertiary alcohol/ halogenolkane
Carbon it is attached to is attached to three carbon atoms
Steric hindrance
Bulky groups make it difficult for an incoming group to attack the carbon atom
3 charge centres
120
Planar triangular
Pi bond
Bond formed by the sideways overlap of parallel p orbitals
Isomer
Same molecular formula different structural and geometric formula
First ionisation energy
Minimum energy needed to remove one mole of electrons from one mole of gaseous atoms in their ground state
Metallic bonding
Sea of delocalised electrons
Lattice of positive ions
Electrostatic attraction
Exothermic
A reaction that releases heat to the surroundings as a results of forming products with stronger bonds than the reactants.
Endothermic
A reaction that absorbs heat from the surroundings as a results of forming products with weaker bonds than the reactants.
Standard enthalpy change of reaction
Heat energy transferred during a reaction under standard conditions of temperature (298k) and pressure (1atm)
Average bond enthalpy
Energy needed to break one mole of bonds in gaseous molecules under standard conditions (298k
Standard state
Pure form of the substance under standard conditions of 298k
Standard enthalpy of formation
Enthalpy change that occurs when one mole of the substance is formed from its elements in their standard states under standard conditions of 298k and
100kpa
Products – reactants
Standard enthalpy change of combustion
Enthalpy change when one mole of the substance undergoes complete combustion under standard conditions (298k and 100kpa) and where the reactants and products are in their standard states.
Reactants – products
Hess law
Enthalpy change for any chemical reaction is independent of path provided the starting conditions and final conditions, and reactants and products, are the same.
Charge centres : 2
Linear
180
Homologous series
Differ by CH2
Similar chemical properties
Graduation in physical properties
Solubility of alkane
Non polar therefore do not form hydrogen bonds. Also not polar therefore colloid of water – not soluble.
Volatility of alkanes
Increase molecular size = increased van der walls = increase boiling point = decrease volatility
Lower men = gases
Ammonia v ammonium
Ammonia =NH3
Ammonium = NH4
How do hydrogen bonds and dipole to dipole attraction affect volatility of organic compounds?
Increase strength of bonding therefore decrease volatility
Volatility of functional groups. Highest volatility first.
Alkene – halogenoalkane – aldehyde – ketone – alcohol – carboxylic acid.
Solubility of functional groups.
Most soluble: alcohol, carboxylic acid, amine (due to hydrogen bonds)
Less soluble: aldehydes, ketones, amine, ester
Incomplete combustion
Carbon monoxide / carbon and water.
NEVER HYDROGEN GAS!
Alkane bonds
High bond enthalpies as bonds are non polar
REACTION MECHANISM: between alkane and halogen.
Initiation, termination and propagation steps
Verbally explain reactions mechanism between halogen and alkane
Homolytic fission: splitting shared pair of electrons equally between two products.
Uv light
Then initiation, propagation and termination
Compete combustion of alcohol produces
Carbon dioxide and water
Describe using equations the oxidation reactions of primary alcohols. Conditions.
1. Acidified potassium
Dichromate.
TWO STEP REACTION
2. Alcohol – aldehyde – carboxylic acid.
Oxidising agent for oxidisation of alcohols.
Acidified potassium dichromate
=
(O)
How do you obtain aldehyde as product? And why is it possible?
Distill it as it forms. Possible because aldehydes have lower boiling points.
How do you obtain carboxylic acid in oxidation of primary alcohol?
Want aldehyde to remain in contact with oxidising agent for as long as possible. Therefore heat under reflux using reflux condenser.
Are tertiary alcohols oxidised by potassium dichromate?
We are going to assume not.
Products formed by oxidation of primary and secondary alcohols.
Primary alcohol – aldehyde – carboxylic acid

Secondary alcohol – ketone ( under reflux)

Alkane 2 halogenoalkane
– use halogen
– substitution reaction
– uv light
RECALL INITIATION, prop
And term
Alkenes to alkane
– hydrogen
– hydrogenation
– nickel catalyst
– 150
Alkene to dihalogeno alkane
– use halogen
– room temp
– loss of colour
Alkene to halogenoalkane.
Use halide
Room temp
In solution
In Alkene to halogenoalkane which halide reacts most readily- why?
HI – weakest bond therefore reacts most readily.
Alkene to alcohol
Water
Hydration
High pressure with steam
Concentrate sulphuric acid.
Polymerisation – type of reaction and conditions
– addition reaction
– high pressure
– high temperature
– catalyst
Economic importance of alkenes
Alkenes used in addition polymerisation to make synthetic plastics and cloth.
Large industry
Alcohol combustion products
Carbon dioxide and water
Alcohol to aldehyde
– primary alcohol
– oxidising agent = acidified potassium dichromate
– distilled
Alcohol to carboxylic acid
– primary alcohol
– oxidising agent = acidified potassium dichromate
– aldehyde to remain in contact with the oxidising agent for as long as possible therefore
– heat under reflux.
Halogenoalkanes undergo
NUCLEOPHILIC SUBSTITUTION
AND
ELIMINATION REACTION.
TWO REACTION MECHANISMS: Halogenoalkane to alcohol
– can be SN2 or SN1 depends upon whether primary or tertiary halogenoalkane
– sodium hydroxide
– SN2 – two reactants on one – transition state – rate determined by both
– SN1: Steric
Hindrance prevents transition state. Carbocation instead.
Warm aqueous
Solution
Rate of Nucleophilic substitution
Tertiary = fastest
Primary = slowest
Expect flouroalkane to be most as flouro most electeonegative BUT
HI is weakest bond – need to break this bond therefore HI fastest.
Benzene
6 carbon atoms in a ring
C6H6
Halogenoalkane to alcohol
Warm reactants in aqueous solution.
Primary halogenoalkane = primary alcohol and so forth
REACTION MECHANISM SN2 only:Halogenoalkane to amine
Ammonia
Concentrated ammonia solution
High pressure.
Ammonia reacts to form other products therefore increase concentration to decrease other products.
REACTION MECHANISM sn2: Halogenoalkane to
Nitrile
Substitution reaction
Cyanide ion from potassium cyanide.
Heat under reflux in solution of potassium
Cyanide in ethanol. Alcohol = Solent for polar and non
Polar.
Nitrile to amine
Hydrogen
Nickel catalyst
Product = primary amine.
Elimination =
Introduces unsaturation into molecule.
Remove small molecule from larger molecule.
Halogenoalkane to Alkene
DRAW BOTH
Can be E1
What is a condensation reaction? And what must be present in order for it to occur?
A condensation reaction is when two molecules react to form a product with the loss of a small molecule.
In order for it to occur: must have two functional groups on each side.
Alcohols to esters
Carboxylic acid and alcohol = ester and water
Condensation reaction
Warm carboxylic acid and h2so4 (catalyst)
Uses of esters
Sweet fruity smells
Food flavourings
Perfumes
Plasticisers
Naturally occurring fats and oils.
Solvents
Amine carboxylic acid =
Amides
COOH NH2 = CONH(R) H2O
Deduce structure made from 1,6 – diamino hexane and hexanedioic acid. Identify repeating unit. How is it
Formed?
Polyamide making. Amine carboxylic acid = Amide.
Economic importance of condensation reactions.
Make synthetic polymers – large industry. No need for cotton and silk
Describe Geometrical isomerism BOTH CASES!
Alkene: double bond in alkenes causes restricted rotation as free rotation would cause the pi bond to break.
Cycloalkanes: the ring in cycloalkanes causes restricted rotation due to strained bond angles.

A result of restricted rotation (must mention phrase).
CIs = xx, yy
Trans = xy, yx

Physical and chemical properties of CIs isomers
Higher coiling point than Trans as has dipole
Lower boiling point than Trans as
Cannot pack together as well.
Fairly similar chemical properties
Except
CIs and Trans -1,2-dichloroethene
Boiling point diff
CIs and tran but-2-ene-1,4-dioc acid. Explain difference in bonding and therefore difference in melting points and solubility. What happens when heated
Cis: forma intramolecular hydrogen bonds (draw and you will see) therefore lower melting and boiling point. Less soluble.
Trans: intermolecular hydrogen bonds. Higher melting boiling point. More soluble.
Sublimes = Trans.
CIs = anhydride
Explain geometrical isomerism in cycloalkanes
Ring prevents rotation. Therefore can be CIs and Trans.
Optical isomerism
Asymmetric or chiral carbon = 4 diff groups.
Enantiomers
Two superimposible forms of an
Optical isomer.
Racemic mixture
A mixture containing two equal amounts of the two Enantiomers. Is not optically active
Physical and chemical properties of the Enantiomers .
Identical except for optical activity: rotate plane of polarisation in equal and opposite directions.
Reactivity with other chiral molecules: React to produce two very different and distinct products. Two Enantiomers react same with non chiral molecules.
Optical isomers
Rotate the plane of polarisation.
Uses of radioisotopes. Give the THREE EXAMPLES!
Carbon 14: dating – relative abundance is constant. Carbon 14 is always present in living things. Upon death don’t absorb carbon 14. Level of carbon 14 falls – use of half life can give estimate of relative date of object.
Cobalt 60: radiotherapy. Treat cancer with ionising radiation. Knocks off electrons from cells = damaged. normal recover cancerous cannot recover. Penetrating gamma radiation
Iodine131: emits beta and gamma. Thyroid cancer. Iodine 125: prostate cancer. Medical tracer – detect radiation levels.
Radioisotopes chemical properties compared to parent atom
Isotopes of elements that undergo radioactive decay. Radioisotopes share same chemical properties as atoms and have same role in body as their parent element.
Relative atomic mass
AVERAGE mass of an atom of the element taking into account all it’s isotopes and their relative abundance, compared to one atom of carbon 12.
Continuous v line spectrum
Continuous = mixture of all wavelengths
Line: specific wavelength
How are lines in H emission spectrum related to energy levels?
Energy of photo = hf therefore different energy level transitions correspond to different wavelength of light.
8-3 = infra
7-2 = visible
6-1= ultraviolet
Thinking about electrons in ground state – lowest energy level =
Nearest to nucleus as least potential energy
The lines in the visible emission spectrum of hydrogen atoms converge at
Higher energy levels = 8 = furthest away from nucleus
Coordination number
Number of ions that surround a given ion in the lattice
Measuring the PH
Involves
Measuring the hydrogen ion concentration. Dont measure electron transfer regardless of lewis theory
Outline characteristics of state of
Equilibrium
Concentration does not change.
Dynamic: forward and backward reaction happen at the same rate.
Closed system: no exchange of matter with surroundings.
No change in macroscopic properties
Strong base examples (LBNK)
LiOH
Ba(OH)2
NaOH
KOH
Weak bases
Ammonia (NH3) ethylamine (C2H5NH2)
Strong acids
HCL
HNO3
H2SO4
Weak acids
Ethanoic acid (CH3COOH)
Carbonic acid (H2CO3)
Phosphoric acid (H3PO4)
Kw
Concentration of water =
Constant at specific temp therefore
Combine with kc = Kc[H2O] = Kw
Kb
Base disassociation constant
Kc[H2O]
Electronegativity
Ability of an atom to attract electrons in a covalent bond.
Atomic radii Trends
– group
– period 3
1. Atomic radii increases as no of occupied shells increases.
2. Decrease as increase in CHARGE DENSITY (atomic no) greater attraction
Ionic radii trends
– group
– period 3
1. Increases down a
Group
2. Across a period = decrease till group 4 ( increase in nuclear charge)
= 4:7 decrease due to increase in nuclear charge
3. Postive ions are smaller
Than negative ions
Melting point
– alkaline metals
– halogens
– decrease down group 1 increased distance between lattice and electrons
– increase down group 7 = increase in no of electrons = stronger van der waals.
First ionisation energy trends
– period
– group
Increase across period. -Increase atomic no.
Decrease down a group – effective nuclear charge remains constant while increased shielding from inner electrons.
Provides evidence of sub levels
Electronegativity
– period
– group
Increases across a period
Decreases down a group
Sn1
Unimolecular Nucleophilic substitution
Sn2
Bimolecular Nucleophilic substitution
Characteristic properties of transition metals
Variable oxidation number, complex ion formation, existence of coloured compounds, catalytic properties , partially filled d sub shells
Why are scandium and zn not considered to be transition elements?
Scandium – does not have partially filled d
Sub shell as an ion.
Zinc : does not have partially filled d subshell
Explain why there are variable oxidation numbers in transition metal IONS!
Oxidation states of Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn
Oxidation states of Ti , V, Cr, Mn, Fe, Co, Ni, Cu, Zn
Ligand
Species that uses a lone pair of electrons to form a dative covalent bond with a metal ion.
Why are some complexes of d blocks coloured.
Electric field due to lone pair of electrons splits d sub shell.
Different energy levels. Electrons absorb energy and moves between energy LEVELS.
Frequency of light absorbed means complementary colour is transmitted.
State examples of the catalytic action of transition elements and their compounds. 4 examples requires
MnO2 in decomp of hydrogen peroxide
V2O5 in contact process
Fe in haber process
Ni in conversion of Alkenes to alkanes
Co in via b
Pd and pt in catalytic converters
Economic significance of catalysts in contact and haber process
Reduce costs
Chemical and
Physical properties of nobles
Mono atomic
Unreactive
Colourless
Gases
Chemical and physical properties of group 1
Highly reactive
Good conductors of electricity
Low density
Grey shiny surfaces
Form ionic compounds
Chemical and physical properties of group 7
Coloured
Gradual change from gases to liquid
Very reactive
Reactivity decreases down the group
Form ionic compounds
Or covalents
Ionic bond
Electrostatic attraction between oppositely charged ions.
Nitrate
NO3-
Hydroxide
OH-
Hydrogen carbonate
HCO3-
Carbonate
CO3)2-
Sulfate
SO4)2-
Phosphate
PO4)3-
Ammonium
NH4)
Covalent bond
Electrostatic attraction between a pair of electrons and positively charged nuclei
3 charge centres
1 lone pair
V shaped
117.5
4 charge centres
0 lone pair
Tetrahedral
109.5
4 charge centres
1 lone pair
Pyramidal
107
4 charge centres
2 lone pair
V shapes
104.5
5 charge centres
0 lone pair
Triangular bipyramidal
180, 120, 90
5 charge centres
1 lone pair
See saw
117, 90
5 charge centres
2 lone pair
T shapes
90
5 charge centres
3 lone pair
Linear
180
6 charge centres
0 lone pair
Octahedral
90
6 charge centres
1 lone pair
Square pyramidal
6 charge centres
2 lone pair
Square planar
90
Diamond structure
And bonding
1 carbon attached to four other carbon atoms
Sp3
Tetrahedral
C60 – fullerene mention 4 things.
Sp2 hybridised
60 atoms in a sphere
Accepts electrons to form negative ions – semiconductor
1 carbon bonded to three other carbons
Graphite
1 carbon to three other carbon atoms
Delocalised electrons
Conducts
Sp2
Silicon
One silicon atom to four other silicon atoms. Tetrahedral arrangement.
Giant covalent.
Silicon dioxide SiO2
Giant covalent structure
Each si is bonded to four oxygen atoms
Each O bonded to two
Si atoms
Hybridisation
Mixing of atomic orbitals to form new orbitals for bonding
Rate of reaction
Change in concentration of reactants per unit time
Activation energy
Minimum kinetic energy particles must have before they are able to react
Rate of reaction depends on: 6
Temperature
Concentration
Particle size
Pressure
Catalyst
Geometry of atoms
Maxwell Boltzmann energy distribution curve.
Number of particles with kinetic energy = y
Kinetic energy
= x
Does temperature affect the rate constant?
Yes an increase in temperature increases the rate constant.
What MUST you remember when doing buffer calculations??
1. Disassociation of weak acid is so small = 0
Therefore concentration at eq = initial
2. Salt is fully disassociated
Standard states of period 3 elements
Na(s)
Mg
Al
Si
P4
S
Cl2(g)
Period 3 oxides and standard states
Na : Na2O (solid)
Mg: MgO (solid)
Al: 2Al2O3 (solid)
Si: SiO2 (solid)
P: P4O10 P4O6 (solid)
S: SO2 (solid) SO3 (liquid)
Cl: Cl2O7 (g) Cl2O (g)
Electrical conductivity and bonding of period 3 oxides in molten state
Na: high, Giant ionic
Mg: high, Giant ionic
Al2O3: high, Giant ionic
SiO2: very low, giant covalent
P4O10/P4O6: none, molecular covalent
SO3/SO2: none, molecular
covalent
Cl2O7/ Cl2O: none, molecular covalent
Period 3 oxides
Water ( only need to do Na2O, MgO, P4O10, SO3). Do Al2O3 and SiO2 react? Therefore period acid/base nature
NaOH aq
Mg(OH)2 aq
H3PO4 (aq)
H2SO4 (aq)
No they do not react with water.
Period 3 chlorides and their standard states
Na : NaCl solid
Mg: MgCl2 solid
Al: AlCl3 solid AlCl6 gas
Si: SiCl4 liquid
P: PCl3 liquidPCl5 solid
S: S2Cl2 liquid
Cl: Cl2 gas
Period 3 chlorides and chlorine
water. Therefore determine acid base nature
Chlorine
Water ~ HCl HOCl
NaCl: hydrated into ions – neutral
MgCl2: hydrated into ions – weakly acid
AlCl3: Al(OH)3 3HCl
SiCl4:SiO2 4HCl
PCl3
Difference between AlCl3 and Al2Cl6
AlCl3 = solid
Al2Cl6 = liquid. Dimer because of two dative covalent bonds from Cl to al
Electrical conductivity of period 3 chlorides and bonding
Na: highly conductive
Mg: highly conductive
Al: poor conductor
Rest = no conductivity
Li, Na, K water
LiOH
NaOH
KOH
Explain how a redox reaction is used to produce electricity in a voltaic cell.
In diagram. Separate a spontaneous redox reaction into two half cells and allow electrons to transfer between by an external circuit.
Essential components of voltaic cell.
Voltmeter
Salt bridge
Two half cells
Essential components of electrolytic cell.
Battery/ cell
Electrolyte
Electrodes
In voltaic cell where does oxidation occur? And where does the current flow from and to?
Anode = oxidation
From cathode to abode
In electrolytic cell where does oxidation take place and where does current flow and to?
Anode = oxidation
From anode to cathode
Ions present in electrolysis of molten salts
Only ions present are from salt.
Writing half equations
Electrons to balance oxidation numbers.
Protons to balance charge
Add h20 to balance O2
Eliminate electrons by making them on opposite sides and equalising the magnitude
Periodicity exam notes:
charge density not just charge
Strong acid weak base titration point is
Ph5
What happens at titration point weak acid strong base? Why is titration point above 7ph?
At titration point all acid is used up. Resulting ph is dependent on poh as only base from partial disassociation of weak acid left. No strong base or partial disassociation of weak acid. Therefore higher than 7 ph
Titration point on graph =
Steepest point
State hasselbalch henderson equation
Ph = pka log [partial disassociated acid or base][HA]
Alcohol to ketone
Secondary alcohol
Acidified potassium dichromate
Heat under reflux
Sec alcohol [O] ~ ketone water
Alcohol to
Aldehyde and then carboxylic acid equations
Alcohol [O] ~ aldehyde water
Aldehyde [O] ~ carboxylic acid water
Equivalence point in strong acid and base
Strong base completely neutralises strong acid.
HCL ~ Cl and h
Cl is not base = neutral therefore neutral ph at 7
State relationship between HA and A- or BOH and B at half equivalence point
Equal concentrations
A buffer is most effective at…
Half equivalence point
High concentration of NaCl in aqueous solution. What is produced at each electrode.
When concentrated = Cl and h2
Low concentration= O2 and h2
How do you electroplate?
Electrolyte of ions to be deposited.
Cathode = metal to be plated.
Anode = inert electrode.
Obs: cathode getting plated, possible loss of colour if Cu2 = blue
Observations alkali metals added to water.
Hydrogen gas evolved.
Temp of water increases.
Clear, colourless solution is formed.
Properties of alkali metals
Soft. Shiny. Silvery.
SiO2 bonding
Four oxygen to one si .
Both oxygen and si are sp3 hybridise. Each oxygen connected to two si
When describing emission spectra must mention
Convergence at higher energy levels
Rate for A
2A
A B
K=rA
K=rA^2
K=r [A][B]
Choosing indicator look at pH at …
Equivalence point.
Enthalpy change of atomisation .
Heat change that occurs when one mole of gaseous atoms are formed from the element in its standard state.
Electron affinity
enthalpy change when one mole of gaseous atoms attracts one mole of electrons.
Second electron affinity
enthalpy change when one mole of gaseous negatively charged ions attract one mole of electrons.
Lattice enthalpy
Enthalpy change that occurs when one mole of a solid ionic compound is separated into gaseous ions under standard conditions.
Kc and rate constant only change for
Temperature. Does not change with pressure, concentration etc
Difference in conditions between Nucleophilic substitution and elimination.
Sub: dilute solution, aqueous, warm
Elim: concentrated, ethanolic, hot
Region of em spectrum used in HNMR spectroscopy.
Radio waves
Provide energy needed to flip nuclei.
Sub level with highest energy
S
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