Definitions

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Definitions

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Energy level

• is the fixed energy value that an electron in an atom may have.

Exam Q (Hons)

‘08/Q10c

‘07/Q4

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Ground State

• lowest energy state ( in 1s orbital)

• Excited state = higher energy state

Exam Q (Hons)

08/Q10c

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An orbital

• is a region in space within which there is a high probability of finding an electron.

Exam Q (Hons)

‘06/Q5

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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.

Exam Q (Hons)

’06/Q10a

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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• relative to 1/12th mass of carbon 12 atom

(expressed on a scale in which the atoms of carbon 12 isotope have a mass of

exactly 12 units).

Exam Q (Hons)

’06/Q10a

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Mass Spec.

• V• I• A• S• D

Victor

A+

Vaporisation

Ionisation

Acceleration

Separation

Detection

how? why?

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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 a measure of the relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond.

Exam Q (Hons)

’06/Q5

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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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Intermolecular Forces

• attractive (repulsive) forces between molecules

• Intramolecular forces are

attractive (repulsive) forces within a molecule

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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

• Getwhite 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

Equation needed !!

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CO32- /HCO3

-

• Adddil. HCl (or any acid)

• Get bubbles of CO2 (limewater milky)

• Distinguish• add

MgSO4 to fresh solution• Get

white ppt. immediately = carbonatewhite 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

• Addammonium molybdate

• Add 5 drops of conc. nitric acid (warm the solution)

• Get yellow ppt

No equation neededConfirm: Goes colourless when add dilute NH3

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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.

Exam Q (Hons)

07/Q4

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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 one mole of 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

Exam Q (Hons)

‘07/Q10b

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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 gasesunder all conditions of temperature and pressure.

• (Know the assumptions)

Exam Q (Hons)

’06/Q11a

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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

Exam Q (Hons)

’07/Q7

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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

Exam Q (Hons)

‘07/Q7

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Primary Standard

• is a substance of high Mr 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

oxidation number

Exam Q (Hons)

‘08/Q10(b)

’06/Q10(b)

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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 I- to I2

» (known concentration) (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? )

• ClO- oxidises I- to I2

• I2 v. thiosulfate

• starch indicator as before

NB

remember dilution factor

in calculations

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Definitions

Rate of reactions

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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.

Exam Q (Hons)

2003 Q7

2004 Q8

2007/Q9

2011 /Q5

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Factors affecting rate

• nature of reactants

• particle size

• concentration

• temperature

• catalysts

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Equations to know

• Write each equation then check • Decomposition of hydrogen peroxide using

manganese dioxide as catalystMnO2

• 2H2O2 2H2O + O2

• Sodium thiosulfate and hydrochloric acid• Na2S2O7 + 2HCl S + 2NaCl + SO2 +H2O

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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.

Exam Q (Hons)

2003 Q7

‘07/Q9

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Homogeneous catalysis

• occurs when the reactants and the catalyst are in the same phase.

• example =?liquids

• KI catalyses 2H2O2 2H2O + O2 (iodine snake)

• And any enzyme

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Heterogeneous catalysis

• occurs when the reactants and the catalyst are in different phases.

(NB must be phases not states)

• example = ?• Al2O3 (solid)catalyses

• ethanol (gas) ethene

Exam Q (Hons)

‘07/Q4

Methanol methanal using platinum

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Autocatalysis

• occurs when one of the products of the reaction catalyses the reaction.

• Example = ?• Mn2+ ions in KMnO4 titrations

(purple changes to colourless more quickly as titration proceeds)

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Mechanism of Catalysis

• Intermediate Formation theory

• Surface Adsorption theory

• Know details of each and evidence of intermediate formation theory

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Enzymes

• Are biological catalysts made of protein • Examples of homogeneous catalysis• Need to know 2 examples

– Amylase catalyses conversion of starch to maltose– Catalase catalyses conversion of hydrogen peroxide to

hydrogen and water

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Learning Check

• Do I know Definition for

1. Rate of reaction2. Catalyst 3. Homogeneous catalysis 4. Heterogeneous catalysis 5. Auto catalysis 6. Two mechanisms of catalysis

Press enter to continue

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Catalytic converter• Catalysts = ?• Pt + Pd + Rh on honeycomb surface

(ceramic) • Gases in

CO NO

NO2

hydrocarbons • Gases out

CO2 and N2 and H20

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Learning Check

• Do I know

1. 3 metals in Catalytic converter 2. 4 wastes in exhaust fumes 3. Problem of each4. What each is converted to 5. What poisons catalytic converter 6. Type of catalysis occuring in catalytic converter

Press enter to continue

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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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Effective Collision

• Is one in which a reaction occurs

• The activation energy has been reached or exceeded.

Exam Q (Hons)

2009 Q9

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Activation Energy

• is the minimum energy which colliding particles must have for a reaction to occur

(minimum energy required for effective collisions between particles)

Exam Q (Hons)

2006/Q7

2009/Q9

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Activation Energy 2

• Catalysts lower the activation energy of a reaction

Without catalyst

with catalyst

Compare E act

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Energy Profile Diagram

• Sketch an energy profile diagram for an endothermic reaction.

• Press enter when ready and• It should look like this

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EndothermicEnergy Profile Diagram

0 2 4 6 8 10 120

1

2

3

4

5

6

7

Energy

Reactants

Products

Energy In

Activation Energy

Note – axes should be labelled Time (x) and energy (y)Curve should be smooth !

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Energy Profile Diagram

• Sketch an energy profile diagram for an exothermic reaction.

• Press enter when ready and• It should look like this

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ExothermicEnergy Profile Diagram

0 2 4 6 8 10 120

1

2

3

4

5

6

7

Energy

Reactants

Products

Energy Out

Activation Energy

Note – axes should be labelled Time (x) and energy (y)Curve should be smooth !

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Learning Check

• Do I know Definition for 1. Effective collision

2. Activation energy

Can I draw energy profile diagram for 1. Exothermic reaction2. Endothermic reaction3. Either of above with catalyst

The End

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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

reactions at equilibrium // oppose the applied stress(es)*

Exam Q (Hons)

’06/Q11b

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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 = ----------------- for aA + bB cC + dD[A]a x [B]b

(product 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+]

• pH < 7 acid• pH = 7 neutral• pH > 7 base

[ ] =

moles per litre

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Kw

• Kw = [H+].[OH-]

([H+] = √Kw)

• Also remember

Kw = 1x10-14 ( at 25oC)

so [H+]= 1x10-7

and pH =7

Exam Q (Hons)

08/Q8

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Strong / Weak acid

• A strong acid is a good proton donor or (is fully dissociated into ions in dilute aqueous soln.

• [H+] = [acid] HCl• [H+] = 2x[acid] H2SO4 etc

• A weak acid is a poor proton donor or (slightly dissociated into ions in dil. aq. soln.)

• [H+] = √Ka x Macid

Exam Q (Hons)

‘07/Q7

AG

Strong / Weak base

• A strong base is a good proton acceptor or one which is fully dissociated into ions in dilute aqueous solution

• [OH-] = [base] NaOH• [OH-] = 2x[base] Ca(OH)2 etc

• A weak base is a poor proton acceptor or one which is slightly dissociated in dli. aq. soln.

• [OH-] = √Kb x Mbase

AG

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.

Exam Q (Hons)

’06/Q8

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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

AG

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 at least one carbon – carbon double or triple bondand undergoes addition reactions

• alkenes / alkynes

• test for unsaturation decolourise bromine solution

Exam Q (Hons)

08/Q9

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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)

Exam Q (Hons)

‘08/Q6

’06/Q6

AG

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 by heat and catalysts

(for which there is a greater demand)

Exam Q (Hons)

‘07/Q6

AG

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)

AG

Exothermic reaction

• is one which produces heat.

• ∆H is minus ( giving away)

AG

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.

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Chromatography

• is a separation technique in which

a mobile phase

carrying a mixture

moves in contact with

a selectively adsorbent stationary phase.

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Instrumentation

• Mass Spec.• AAS• GC• HPLC• IR spec.• UV spec. • X ray crystallography (option 2)

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Mass Spec.

• Positively charged ions are separated

• according to different relative masses

• when moving through magnetic field

Victor

A+

VapourisationIonisation AccelerationSeparationDetection

Used toAnalyse blood of race horses for drugsIdentify substances

Principles Processes

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Atomic Absorption Spectrometry

• 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)

DissolveAtomise AbsorbMeasure Detection

Used to analyse water samples for heavy metals Cd Hg Pb

Principles Processes

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Gas Chromatography GC

• Different components have different tendencies to dissolve in a non-volatile liquid, which is coated on fine particles of a solid in a the GC column

Injection …Transport …Separation …Detection ...

Used with MS in drug testingalso blood alcohol levels

Principles Processes

mobile phase ?

stationary phase??

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HPLC

• High Performance Liquid Chromatography

• Different components of a mixture have different tendencies to adsorb onto fine particles of solid in HPLC column

Principles Processes

Injection …Transport …Separation …Detection …

Used to separate less volatile mixtures e.g. growth promoters in meat.mobile

phase ?stationary phase??

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IR

• Infra red spectrometry • Molecules of a substance

absorb infra –red of different frequencies.

(different number/ type bonds)

• The combination of frequencies absorbed is unique to the molecules of each substance

Prepare …Transmit IRAbsorption… Detection…Spectrum obtained

Used to identify functional groupsand identify drugs

Principles Processes

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UV

• Ultra violet spectrometry

• Molecules absorb UV radiation

• Electrons promoted from ground state to higher energy states.

• Absorption is directly proportional to concentration.

Prepare Transmit UV through

Blank (o%abs)Sample (known + unknown)

Spectrum obtained

Quantatativeused to find amount of org. subs. e.g. drugs

Principles Processes

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X ray crystallography

• 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

Prepare ….Transmit : x-ray detected on filmPattern analysed and structure worked out

Used to determine structure of macro-molecules e.g. DNA

Principles Processes