Download Definitions - Loreto Science

Definitions
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Energy level
• is the fixed energy value that an electron
in an atom may have.
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An orbital
• is a region in space within which there is a
high probability of finding an electron.
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An element
• is a substance that cannot be split up into
simpler substances by chemical means.
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A triad
• is a group of three elements with similar
chemical properties in which the atomic
weight of the middle element is
approximately equal to the average of the
other two.
(Dobereiner)
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Newlands’ Octaves
• are groups of elements arranged in order
of increasing atomic weight, in which the
first and the eighth element of each group
have similar properties.
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Mendeleev’s Periodic Law
• When elements are arranged in order of
increasing atomic weight (relative atomic
mass), the properties of the elements vary
periodically.
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The atomic number(Z)
• is the number of protons in the nucleus of
that atom.
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Periodic Table
• is an arrangement of elements in order of
increasing atomic number.
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Elements are arranged
• in order of increasing atomic number, the
properties of the elements vary
periodically.
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Mass number (A)
• is the
sum of the number of protons and neutrons
in the nucleus of an atom of that element.
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Isotopes
• are atoms of the same element ( i.e. they
have the same atomic number) that have
different mass numbers due to the
different number of neutrons in the
nucleus.
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Relative Atomic Mass
• is the average of the mass numbers of the
isotopes of the element
• as they occur naturally
• taking their abundances into account
• expressed on a scale in which the atoms
of carbon 12 isotope have a mass of
exactly 12 units.
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Mass Spec.
•
•
•
•
•
V
I
A
S
D
Vaporisation
Ionisation
Acceleration
how? why?
Separation
Detection
A+
Victor
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Aufbau Principle
• that when building up the electronic
configuration of an atom in its ground
state, the electrons occupy the lowest
available energy level.
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Hund’s Rule of Maximum
Multiplicity
• states that
when two or more orbitals of equal energy
are available,
the electrons occupy them singly first
before filling them in pairs.
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Pauli Exclusion Principle
• that no more than two electrons may
occupy an orbital and they must have
opposite spins.
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Compound
• is a substance that is made up of two or
more different elements combined
together chemically.
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Octet Rule
• that when bonding occurs, atoms tend to
reach an electron arrangement with eight
electrons in the outermost shell.
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An Ion
• is a charged atom or group of atoms.
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An Ionic bond
• is the force of attraction between
oppositely charged ions in a compound.
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A transition metal
• is one that forms at least one ion with a
partially filled d sublevel.
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Molecule
• is a group of atoms joined together. It is
the smallest particle of an element or
compound that can exist independently.
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Valency of an element
• is defined as the number of atoms of
hydrogen or any other monovalent element
with which each atom of the element
combines.
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Electronegativity
• is the relative attraction that an atom in a
molecule has for the shared pair of electrons
in a covalent bond.
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Electronegativity
• difference > 1.7 indicates ionic bonding in
a compound.
• An electronegativity difference ≤ 1.7
indicates covalent bonding in a compound.
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The value of electronegativity
• decrease down the groups in the Periodic
Table for two reasons:
• increasing atomic radius
• screening effect of inner electrons
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The values of electronegativity
• increase across the periods in the Periodic
Table for two reasons:
• increasing nuclear charge
• decreasing atomic radius
F= most electronegative element.
Halogens –decrease in reducing power down the group
due to drop in electroneg. values.
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Vans der Waals Forces
• are weak attractive forces between
molecules resulting from the formation of
temporary dipoles.
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Dipole-dipole
• Dipole – dipole forces are forces of
attraction between the negative pole of
one molecule and the positive pole of
another.
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Hydrogen bonds
• are particular types of dipole-dipole
attractions between molecules in which
hydrogen atoms are bonded to nitrogen,
oxygen or fluorine.
• The hydrogen atom carries a partial
positive charge and is attracted to the
electronegative atom in another molecule.
Thus, H acts as a bridge between two
electronegative atoms.
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The Law of Conservation of Mass
• the total mass of the products of a
chemical reaction is the same as the total
mass of the reactants.
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The Law of Conservation of Matter
• that in any chemical reaction, matter is
neither created nor destroyed but merely
changes from one form into another.
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Tests for Anions
•
•
•
•
•
Chloride
Sulfate/sulfite
carbonate/hydrogen carbonate
nitrate
phosphate
• (NB know confirmatory test too!)
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Chloride
• Add
AgNO3
• Get
white ppt
• Confirm = ppt dissolves in
dilute ammonia
• Equation needed
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Sulfate/sulfite
• Add
BaCl2
• Get
white ppt
• Distinguish
• add
dil HCl to white ppt
• ppt remains = sulfate
• ppt dissolves = sulfite
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Equation needed !!
CO32- /HCO3• Add
dil. HCl (or any acid)
• Get
bubbles of CO2 (limewater milky)
• Distinguish
• add
MgSO4 to fresh solution
• Get
white ppt. immediately = carbonate
white ppt on heating = hydrogen carbonate
Equation needed !!
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Nitrate
• Brown Ring Test
• Add
fresh FeSO4
At slant add conc. H2SO4 drop wise
• Get
brown ring at junction of 2 layers
No equation needed
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Phosphate
• Add
ammonium molybdate
• Add
5 drops of conc. nitric acid
• Get
yellow ppt (warm the solution)
?Query
confirm’y test
No equation needed
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The atomic radius of an atom
• is defined as half the distance
between the nuclei of two atoms
of the same element that are
joined together by a single covalent bond.
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The values of atomic radius
• increase down any one group in the
Periodic Table for two reasons:
• extra shell
• screening effect of inner electrons
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The values of atomic radius
• decrease from left to right across a
Periodic Table for two reasons:
• increasing nuclear charge
• no increase in screening effect
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The first ionisation energy of an
atom
• is the minimum energy required
to completely remove
the most loosely bound electron
from a neutral gaseous atom
in the ground state.******
•
•
2004 =9 marks (2.25%)
2002 = 8marks(2%)
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The values of ionisation energy
• decrease down the groups in the Periodic
Table for two reasons:
• increasing atomic radius
• screening effect of inner electrons
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The values of ionisation energy
• increase across the Periodic Table for two
reasons:
• increasing nuclear charge
• decreasing atomic radius
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More on ionisation energy
• First Ionisation Energy
• M – e- M+
• Second Ionisation energy
M+ – e- M2+
Major jump in I.E. values – significance
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The value of electronegativity
• decrease down the groups in the Periodic
Table for two reasons:
• increasing atomic radius
• screening effect of inner electrons
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The values of electronegativity
• increase across the periods in the Periodic
Table for two reasons:
• increasing nuclear charge
• decreasing atomic radius
F= most electronegative element.
Halogens –decrease in reducing power down the group
due to drop in electroneg. values.
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A gas
• is a substance that has no well-defined
boundaries but diffuses rapidly to fill any
container in which it is placed.
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Radioactivity
• is the spontaneous breaking up of
unstable nuclei
• with the emission of one or more types of
radiation.
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Alpha particles
•
•
•
•
loss of He nucleus (2p + 2n)
mass number down by 4
atomic number down by 2
element changes to element two places
back
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Beta particle
• neutron changes to proton and electron
• electron emitted
• mass number stays same
• atomic number drops by one
• element changes into element one place
back
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Gamma radiation
• no new atoms formed
• (no transmutation)
• only energy lost
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Half Life
• of an element is
the time taken for half the nuclei
in any given sample to decay.
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Mole
• is the amount of a substance which
contains 6 X 1023 particles of that
substance
• (avogadro’s number or constant =L)
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a few numbers
• Kelvin = Celsius + 273
• standard temp
=
273 K
• standard pressure = 1X105 Pa
(100kPa)
m3 = litres X10 -3
m3 = cm3 X10 -6
(1 litre = 1000cm3)
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Mole
• contains 6 X 1023 particles
• has mass equal to Ar or Mr in grams
• occupies 22.4 litres at s.t.p (if gas)
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Boyle’s Law
• states that:
at constant temperature,
the volume of a fixed mass of gas is
inversely proportional to its pressure.
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Charles’ Law
• states that:
at constant pressure,
the volume of a fixed mass of a gas
is directly proportional to its temperature
measured on the Kelvin scale.
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General Gas Law
• P1 X V1 = P2 X V2
T1
T2
Temp in Kelvin
Units for volume same each side
Units for pressure same each side
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Gay Lussac’s law of Combining
Volumes
• the volumes of the reacting gases and the
volumes of any gaseous products are
in the ratio of small whole numbers
provided the volumes are measured at the
same temp and pressure
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Avogadro’s Law
• states that
• equal volumes of gases contain
• equal numbers of molecules under the
same conditions of temp. and pressure
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Molar Volume
• At s.t.p
• one mole of any gas
• occupies 22.4 litres
• Remember to watch out for r.t.p in
questions
• room temp. and press = as given in Q
• (often 24 litres)
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Ideal Gas
• is one which perfectly obeys
• all the gas laws and all the assumptions*
of the kinetic theory of gases
• under all conditions of temperature and
pressure.
•
(Know the assumptions)
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Real v. ideal gas
• Real gases differ from ideal gases at high
pressure and low temp. because
• there are forces of attraction/repulsion between
the molecules*
• the volume of the molecules is not negligible
compared to the distances between them
(*know examples of real gases and the forces
involved)
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Empirical Formula
• gives the simplest whole number ratio of
the numbers of the different atoms present
in the molecule.
(divide by Ar and get ratio)
(molecular formula is a simple multiple of the
empirical formula)
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Acids / Bases
•
•
•
•
•
Arrhenius Acid + Base
Bronsted Lowry Acid + Base
Neutralisation
Conjugate Acid/ conjugate base
conjugate pair
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Arrhenius Acid and Base
• Arrhenius Acid
• is a substance that dissociates in water to
produce H+ ions.
• Arrhenius Base
• is a substance that dissociates in water to
produce OH- ions.
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Bronsted Lowry Acid /Base
• Bronsted Lowry Acid
• is a proton (H+) donor
• Bronsted Lowry Base
• is a proton (H+) acceptor
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Neutralisation
• is the reaction between
an acid and a base
• forming
a salt and water
(acid + base -> salt + water)
SALT = is formed when the H of an acid is replaced by a metal
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Conjugate Acid / Conjugate Base
• Conjugate Acid
is formed when a base accepts a proton
• Conjugate Base
• is formed when an acid donates a proton.
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Conjugate Pair
• an acid and a base that differ by a proton
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Primary Standard
• is a substance which can be obtained
in a pure stable soluble solid form
so that it can be weighed out and
dissolved in water to give
a solution of accurately known
concentration.
• (Know why high Mr matters)
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Titration
• is a laboratory procedure where a
a measured volume of one solution is added to
a known volume of another solution until the
reaction is complete.
• (concentration of one solution known accurately at start)
• (indicator used to show by colour change when reaction is complete)
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Oxidation Reduction
Revision
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Definitions
• Oxidation is
• addition of
• loss of
• increase in
oxygen
Electrons or hydrogen
oxidation number
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• Reduction is
• loss of oxygen
• gain of electrons
• decrease in oxidation number
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More…
• An oxidising agent causes oxidation
and is itself reduced.
• A reducing agent causes reduction
and is itself oxidised.
• What is a redox reaction?
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What is oxidised and reduced in each of the following?
• Br2 + 2Fe 2+ → 2Br– + 2Fe 3+
• Cu 2+ + Zn  Cu + Zn 2+
• 2Na + Cl2  2NaCl
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Oxidation Number Rules
• The oxidation number of
• an Element is
0
• group One elements is +1
• group Two elements is +2
in
compounds
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The oxidation number of
an
ion is equal to the
charge on the ion
• halogens is
-1
(in binary compounds)
(except ……????)
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• The oxidation number of
H in a compound
is +1
– except
in metal hydrides when it is -1
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• The oxidation number of
O in a compound is
-2
– except (x2)
in peroxides when it is -1 (H2O2)
in OF2 when it is +2 (why?)
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• Oxidation numbers
• add up to zero in a compound
• add up to the charge of a complex ion
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• What is the oxidation number of each
element in :H20
MnO4¯
I2
KBrO3
Na2S2O3
H2O2
NaClO
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KMnO4
•
•
•
•
oxidising agent
purple
read top of meniscus
is reduced from
Mn (VII)  Mn (II) in presence of H+
purple  colourless
• own indicator
(end point = first permanent pink)
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KMnO4
• get brown Mn (IV) if H+ absent
• (which acid MUST be used – why x2)
• not primary standard (x2)
• standardised by titrating against standard
solution of acidified Fe 2+
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H2SO4
• added during KMnO4 titrations to provide
H+ and ensure the complete reduction of
Mn (VII)  Mn (II)
and prevent formation of Mn (IV)
(brown)
• added during prep. of Fe (II) solutions to
prevent oxidation of Fe 2+ to Fe 3+ by
oxygen in the air ( why does this matter?)
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Na2S2O3
• S2O3 2- ion
• reducing agent
• used in photography
• not primary standard – why ?
• standardised by titrating against I 2
• starch indicator – when added and why
• colour change at end point ?
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Iodine I2
• Oxidising agent
• NOT a primary standard (X2)
• Produced when MnO4- oxidises
» (known concentration)
I-
to I2
(in excess)
• Starch indicator – when added? why then?
• Colour change at end point
Blue/black to colourless
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Bleach
• sodium hypochlorite Na+ClO• bleach diluted x10 with distilled water
not de-ionised water (why? )
NB
• ClO- oxidises I- to I2
• I2 v. thiosulfate
remember dilution
factor
in calculations
• starch indicator as before
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Rates of Reactions
• The rate of reaction
is the change in concentration per unit
time of any one reactant or product.
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Factors affecting rate
• nature of reactants
• particle size
• concentration
• temperature
• catalysts
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Rate Graphs
• Concentration v. ( 1 /Time )
• or
• Temp v. ( 1 /Time )
• ( 1 /Time )used as Rate and Time
inversely related
• (shorter time means faster rate)
• be careful with units of 1/time
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Catalyst
• is a substance that alters the rate of
reaction
• but is not consumed in the reaction.
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Homogeneous catalysis
• occurs when the reactants and the catalyst
are in the same phase.
• example =?
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Heterogeneous catalysis
• occurs when the reactants and the catalyst
are in different phases.
• example = ?
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Autocatalysis
• occurs when one of the products of the
reaction catalyses the reaction.
• Example = ?
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Mechanism of Catalysis
• Intermediate Formation theory
• Surface Adsorption theory
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Catalytic converter
• Catalysts = Pt + Pd + Rh
• Gases in
CO
NO
NO2
hydrocarbons
• Gases out
CO2 and N2 and H20
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Collision Theory
• for a reaction to occur the reacting
particles must collide with each other
• a collision only results in a product being
formed if a certain minimum energy is
exceeded (called activation energy)
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Activation Energy
• is the minimum energy which colliding
particles must have for a reaction to occur.
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Chemical Equilibrium
• is a state of dynamic balance where the
rate of the forward reaction equals the rate
of the reverse reaction.
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Le Chatelier’s Principle
• If a stress is applied to a system at
equilibrium
• the system readjusts to oppose the stress
applied
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Le Chatelier’s Principle and Gases
• Le Chatelier’s Principle predicts that
in an all-gaseous reaction
an increase in pressure
will favour the reaction which takes place
with a reduction in volume
•
( towards the side with the smaller number of molecules)
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Equilibrium Constant
• Kc
[ ] means concentration in moles per litre
[ C] c x [D]d
Kc = ----------------[A]a x [B]b
for aA + bB cC + dD
(products of products conc.
over
product of reactants conc.)
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Le Chatelier and Industry
• Ammonia and Haber Process
predict max yield at high press. /low temp
reality = 200 atm and 500o C
• Sulfuric Acid and Contact Process
predict max yield at high press. /low temp
reality = one atm and 450oC
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Kc
• large Kc => equilibrium far to right
(lots of product produced)
• small Kc => equilibrium far to left
(v. little product formed)
• must quote temp.
• units – depend on reaction
• tells us how far not how fast a reaction
occurs
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pH
• pH = -log [H+]
[ ]=
moles per litre
• pH < 7 acid
• pH = 7 neutral
• pH > 7 base
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Kw
• Kw = [H+].[OH-]
([H+] = √Kw)
• Kw = 1x10-14 ( at 25oC)
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Strong / Weak acid
• A strong acid is one which is fully
dissociated in solution
• [H+] = [acid] HCl
• [H+] = 2x[acid]
H2SO4 etc
• A weak acid is one which is not fully
dissociated in solution
• [H+] = √Ka x Macid
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Strong / Weak base
• A strong base is one which is fully
dissociated in solution
• [OH-] = [base]
NaOH
• [OH-] = 2x[base] Ca(OH)2 etc
• A weak base is one which is not fully
dissociated in solution
• [OH-] = √Kb x Mbase
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Indicator
• An acid base indicator is a substance that
changes colour according to the pH of the
solution it is in.
• (equilibrium HIn ↔ H+ + In-)
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Methyl orange
• in acid (lower pH )
red
• in base ( higher pH)
yellow
• range
pH 3-5
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Phenolphthalein
• in acid (lower pH )
colourless
• in base ( higher pH)
pink
• range
pH 8-11
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Litmus
• in acid (lower pH )
red
• in base ( higher pH)
blue
• range
pH 5-8
(Not as reliable as others for accurate work)
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Which indicator ?
• strong acid/strong base =
methyl orange / phenolphthalein /litmus( see above)
• strong acid / weak base =
methyl orange
• weak acid /strong base =
phenolphthalein
• weak acid / weak base =
none (why?)
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Hard Water
• is water that will not easily form a lather
with soap
• due to the presence of Ca 2+ or Mg 2+ ions
in solution.
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Temporary Hardness
•
•
•
•
can be removed by boiling the water
due to Ca(HCO3)2
becomes CaCO3 on heating
leads to blocked pipes etc
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Permanent Hardness
• is not removed by boiling the water
• caused by CaSO4 or MgSO4
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Methods of removing hardness
• boiling (only works for temp. hardness)
• distillation
• washing soda
• ion exchange
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Calculations
• Total hardness = calcium hardness +
magnesium hardness
but
• Do calculations as if all hardness caused
by CaCO3
• expressed in p.p.m of CaCO3
• p.p.m. = mg/litre
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Water Treatment
•
•
•
•
•
•
•
screening
flocculation
sedimentation
filtration
chlorination
fluoridation
pH adjustment
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B.O.D
• Biochemical Oxygen Demand is
the amount of dissolved oxygen
consumed by biological action
when a sample of water is kept
at 20oC
in the dark
for five days.
(know reason for each of 3 conditions)
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Eutrophication
• is the enrichment of water with nutrients
which leads to the excessive growth of
algae.
• Nutrients – phosphates/nitrates
• Algal bloom / oxygen depletion
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Sewage Treatment
• Primary Treatment
physical
• Secondary Treatment
biological
• Tertiary Treatment
chemical
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Water Analysis
• Atomic Absorption Spectrometry used to
detect heavy metals like Cd, Hg, Pb
• pH meter
• colorimetry
(Hach test Chlorine in pool water)
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Electrolysis
• is the use of electricity to bring about a
chemical reaction.
• KI/ Acidified water/ Na2SO4/CuSO4 / ions
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Electrolyte
• is a substance that conducts electricity as
a result of the presence of ions.
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Electroplating
• is the process where electrolysis is used to
put a layer of one metal on the surface of
another.
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Electrochemical Series
• is a list of the elements in order of their
standard electrode potentials.
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Organic Chemistry
• is the study of compounds of carbon…
(except some simple compounds
like CO2, CO and carbonates)
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Hydrocarbon
• is a compound that contains
only carbon and hydrogen
• includes
alkanes, alkenes, alkynes
• excludes
alcohols, aldehydes, ketones, carboxylic
acids, esters
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Saturated compound
• is one with only carbon – carbon single
bonds
• alkanes
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Unsaturated compound
• is one which contains carbon – carbon
double or triple bonds
• alkenes / alkynes
• test for unsaturation
decolourise bromine solution
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Homologous Series
• is a series of chemical compounds of
uniform chemical type
• showing gradations in physical properties
• having a general formula for its members
• each member has similar method of prep.
and
• each member differs by (CH2) from
previous member
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Structural isomers
• are compounds with the
same molecular formula but
different structural formulas.
• e.g. butane and methyl propane are both
C4H10
• need to know isomers up to C5H12
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Aliphatic
• an aliphatic compound is an organic
compound that consists of
straight (open) chains of carbon atoms
and closed chain compounds with similar
properties.
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Aromatic
• An aromatic compound is an organic
compound that contains a benzene ring
structure in their molecules.
(benzene – delocalised double bond)
(disc. by Michael Faraday )
(structure by Kekule)
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Octane Number
• of a fuel is a measure of the tendency of
the fuel to resist knocking.
(Best fuels = high octane number
= 100 = 2,2,4 tri methyl pentane )
(Short chains, more branched chains, ring structures)
(Worst fuels = low octane number
=0 = heptane)
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Ways to increase octane number
• isomerisation
• catalytic cracking
• dehydro-cyclis-ation (re-forming)
• add oxygenates
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Isomerisation
• changing straight chain alkanes into
branched chain alkanes
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Catalytic cracking
• is the breaking down of
long chain hydro- carbon molecules into
short chain molecules
(for which there is a greater demand)
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Dehydrocyclisation ( Re-forming)
• involves the use of catalysts to form ring
structures
• straight chain alkanes changed to
cycloalkanes
• cycloalkanes changed to aromatic
compounds
• petrol contains benzene = carcinogen
• health concerns
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Adding Oxygenates
addition of
• methanol
• ethanol
• MTBE
to petrol to increase the octane number.
(Methyl Tertiary Butyl Ether or
2 methoxy 2 methyl propane)
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Exothermic reaction
• is one which produces heat.
• ∆H is minus ( giving away)
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Endothermic reaction
• is one which takes in heat.
• ∆H is positive (add in)
• ammonium nitrate dissolving in water
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Heat of Reaction
• is the heat change involved when the
numbers of moles of reactants indicated in
the balanced equation for the reaction
react completely.
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Heat of Combustion
• is the heat change involved when
one mole of a substance is
completely burned in
excess oxygen
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Kilogram Calorific Value
• of a fuel is the heat energy produced when
1 kg of a fuel is completely burned in
oxygen.
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Bond Energy
• is the energy required to break one mole
of covalent bonds and to separate the
neutral atoms completely from each other.
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Heat of Neutralisation
• is the heat change involved when
one mole of H+ ions from an acid
reacts with
one mole of OH- from a base
forming one mole of H2O
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Heat liberated
• Heat liberated =
M x C x Rise in temp.
Kg
kelvin
• M=Mass of solution in Kg
• c=specific heat capacity
• rise in temp in Kelvin
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Heat of formation
• of a compound is the heat change
involved
when one mole of a compound
in its standard state
is formed from its elements
in their standard states.
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Hess’s Law
• states that if a chemical reaction takes
place in a number of stages,
the sum of the heat changes in the
separate stages is equal to the heat
change if the reaction is carried out in one
stage.
(overall heat change is independent of the
pathway)
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Law of Conservation of Energy
• states that
energy cannot be created or destroyed
but
can be changed from one form of energy
to another.
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Functional Group
• is an atom or group of atoms which is
responsible for the characteristic
properties of a series of organic
compounds.
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Substitution Reaction
• is a chemical reaction in which an atom or
group of atoms in a molecule is replaced
by another atom or group of atoms
• mechanism = free radical substitution
initiation (homolytic fission)
propagation
termination
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Addition Reaction
• is a chemical reaction is which two
substances react together forming a single
substance.
• Mechanism = Ionic addition
• Approach/ polarisation / heterolytic fission
/carbonium ion / product formation
• only happens to unsaturated compounds
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Polymers
• are long chain molecules made by joining
together many small molecules called
monomers.
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Elimination reaction
• is one in which a small molecule is
removed from a larger molecule to leave a
double bond in the larger molecule.
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Organic Synthesis
• is the process of making organic
compounds from simpler starting
materials.
AG
Chromatography
• is a separation technique in which
a mobile phase
carrying a mixture
moves in contact with
a selectively adsorbent stationary phase.
AG
Instrumentation
•
•
•
•
•
•
•
Mass Spec.
AAS
GC
HPLC
IR spec.
UV spec.
X ray crystallography (option 2)
AG
Principles
Mass Spec.
Processes
• Positively charged
ions are separated
• according to
different relative
masses
• when moving
through magnetic
field
Vapourisation
Ionisation
Acceleration
Separation
Detection
A+
Victor
Used to
Analyse blood of race
horses for drugs
Identify substances
AG
Atomic Absorption Spectrometer
Principles
• Ground state atoms of an
element absorb light
characteristic of that
element.
• Absorption is directly
proportional to
concentration.
• (higher absorbance
means higher
concentration of THAT
ELEMENT present)
Processes
Dissolve
Atomise
Absorb
Measure
Detection
Used to
analyse water
samples for heavy
metals Cd Hg Pb
AG
Gas Chromatography GC
Principles
Processes
• Different components
have different
tendencies to
dissolve in a nonvolatile liquid, which
is coated on fine
particles of a solid in
a the GC column
mobile
phase ?
?
stationary
phase?
Injection …
Transport …
Separation …
Detection ...
Used with MS in
drug testing
also blood alcohol
levels
AG
HPLC
Principles
Processes
• High Performance Liquid
Chromatography
• Different components
of a mixture have
different tendencies to
adsorb onto fine
particles of solid in
HPLC column
mobile
phase ?
?
stationary
phase?
Injection …
Transport …
Separation …
Detection …
Used to separate
less volatile
mixtures e.g.
growth promoters
in meat.
AG
IR
Principles
• Infra red spectrometry
• Molecules of a substance
absorb infra –red of
different frequencies.
(different number/ type bonds)
• The combination of
frequencies absorbed is
peculiar to the molecules
of each substance
Processes
Prepare …
Transmit IR
Absorption…
Detection…
Spectrum obtained
Used to identify
functional groups
and identify drugs
AG
UV
Principles
• Ultra violet
spectrometry
• Molecules absorb UV
radiation
• Electrons promoted
from ground state to
higher energy states.
• Absorption is directly
proportional to
concentration.
Processes
Prepare
Transmit UV through
Blank (o%abs)
Sample
(known + unknown)
Spectrum obtained
Quantatative
used to find
amount of org.
subs. e.g. drugs
AG
X ray crystallography
Principles
• Wavelengths of Xrays are
comparable to distance
between atoms in a crystal
• Xrays are scattered when
they hit a crystal surface
• Pattern detected is
analysed and structure
worked out
Processes
Prepare ….
Transmit : x-ray
detected on film
Pattern analysed and
structure worked out
Used to determine
structure of macromolecules e.g. DNA
AG