Ib revision

IB CHEMISTRY ΗL

Revision www.chem.gr Dr. D. Bampilis

IB Chemistry Revision –Dr. D. Bampilis

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IB Chemistry Revision –Dr. D. Bampilis

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IB CHEMISTRY ΗL

Revision

Dr. D. Bampilis

IB Chemistry Revision –Dr. D. Bampilis

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IB Chemistry Revision –Dr. D. Bampilis

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Atomic structure Atomic number (Z): Number of protons in the nucleus.

Mass number (A): Number of nucleons (protons and neutrons) in the nucleus Mass spectrometer very low pressure prevents collisions/unintentional deflections / avoid false readings due to presence of other particles; bombarded with (high energy) electrons to form positive ions;

Electric field = accelerate Magnetic field = deflect Deflection: m/z ↓ more deflection

ions/particles accelerated by electric field/ towards negative plate/cathode; the ions are deflected by the (external) magnetic field;

Uncharged particles are not deflected

factors that affect the degree of deflection: charge (on the ion); higher charge more deflection; mass (of the ion) lighter ions are deflected more than the heavier ions; deflection depends on mass to charge ratio of ions strength of the magnetic field; velocity/speed (of the ions) / strength of electric field; the ions are detected by conversion into an electrical current / the ratio of the intensity of the peaks in the spectrum is equal to the ratio of the ions in the sample Relative atomic mass, Ar ratio of average mass of an atom to 1/12 the mass of C-12 isotope / average mass of an atom on a scale where one atom of C-12 has a mass of 12 / sum of the weighted average mass of isotopes of an element compared to C-12 IB Chemistry Revision –Dr. D. Bampilis

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Calculating relative atomic mass Ar

x1A1 100

(100 x1 )A2 100

Relative molecular mass the average mass of a molecule; compared to 1/ 12 of (the mass of) one atom of 12C / compared to C–12 taken as 12; Isotopes of an element. atoms of the same element/with the same number of protons/with same atomic number but different number of neutrons/mass number/mass; A radioactive isotope of cobalt and its uses. cobalt-60 and radiotherapy/ treatment of cancer, emits gamma radiation which kills cancer cells. sterilization of medical supplies treatment of food sterilizations industrial radiography radioactive tracer cobalt-57 and medical tests/label for vitamin B12 uptake; One use of iodine-131 in medicine and explain why it is potentially dangerous used to investigate functioning of thyroid gland / to treat thyroid cancer / to treat hyperthyroidism; produces gamma rays / destroys healthy cells The use of carbon-14 in carbondating. living organisms have 12C:14C ratio constant/same as atmosphere; after death no more 14C is absorbed and 14C level drops / 12C:14C ratio changes with time / 14C decays / remains become less radioactive; rate of decay of 14C is constant / half-life of 14C is known; measuring radioactivity indicates length of time since death 125

1

I: medical tracer, prostate, brain

H,13C:NMR

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Atomic emission spectra: The characteristic line spectrum that occurs as a result of energy being released by individual elements. Colored lines on a black background.

each transition line is related to energy difference / ∆ E =hf; energy levels are converge

When energy is supplied to an atom, e- are excited from their ground state to a higher energy level.

The e- dropping from higher energy levels to lower emit energy, which can be observed in a spectrum.

As e- can only exist in fixed energy levels, the energy in the emissions are characteristic for each type of atom.

Lines converge toward high energy end of spectrum as the energy levels themselves are convergent.

Atomic absorption spectra: The characteristic line spectrum that occurs as a result of energy being absorbed by individual elements. Black lines on a continuum (colored) background.

at higher energy / f↑ λ↓ c = λ·f Lyman series – Emax - f↑ λ↓– Ultraviolet: from n=2, 3, 4, 5… to n=1 Balmer series – Visible: from n=3, 4, 5, 6, 7 … to n=2 Paschen series – Emin - f↓ λ↑ – Infrared: from n=4, 5, 6, 7 … to n=3

Continuous Spectrum: Shows an unbroken sequence of frequencies (eg. the spectrum of visible light) How a line spectrum differs from a continuous spectrum. continuous spectrum has all colours / wavelengths / frequencies whereas line spectrum has only (lines of) sharp / discrete / specific colours / wavelengths / frequencies;

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First ionization energy energy (per mole) needed to remove one/first/most loosely bound electron from a (neutral) atom; in the gaseous state Ionization energies difference between 2-3 group p orbitals have a slightly higher energy than the 2s orbital/ slightly more distant from the nucleus Ionization energies difference between 5-6 groups The repulsion between the two electrons in the same orbital means that the electron is easier to remove – greater repulsion - Hund’s rule Explain the general increase in successive ionization energies of the element successive electrons (are more difficult to remove because each is) taken from more positively charged ion/ increased electrostatic attraction; Successive Ionization Energies of an element increase: harder to remove an electron from an ion with increasing positive charge

Evidence of shells and subshells o o o o

Each “flat” set gets slighty steeper each time The “steps” get slightly bigger each time First “out” the electrons of outer shell Large increase indicate changes in main energy levels

Electron configuration – Periodic table Half – fill and completely fill subshell d 3d - 4s: Aufbau principle (lowest energy levels are filled first) Cations configuration Hund’s rule: p,d,f subshell- Orbitals within the same sub-shell are filled singly first.

Pauli’s exclusion principle: electrons in single orbital must have opposite spin.

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Bonding Ionic bonding: metal - non metal (ΔΕ>1.5) electrostatic attraction between oppositely charged ions/cations and anions Metalic bonding: metals (conductivity – malleability – ductility - melting point) electrostatic attraction between a lattice of positive ions/cations and delocalized electrons; metals have delocalized electrons which are mobile atoms/ions/layers (of positive ions) can slide over each other without change in the bonding forces higher melting point : greater charge on ion – lower ionic radius Covalent bonding: non metal – non metal (ΔΕ<1.5/ both elements with high Ε) Al2O3: ionic- ΔE=2.0,

Al2Cl6: covalent – ΔE =1.5

Lewis structure showing all valence shell pairs as lines or pairs of dots or crosses Methods for working out Lewis structure -Total number of valence electrons -Join all the atoms together with simple bonds -Arrange the electron pairs so that all the outer atoms have full outer shells -Use the remaining electrons to complete the outer shell of the central atom.

-If there is not electron move lone pairs from outer atoms to become bonding pairs of electrons.

NO2(unpaired electrons) so form N2O4 Strength of the bonds Increasing strength - decreasing length of bond – greatest attractions between the electrons and nuclei -single<double<triple -I<Br<Cl<F bonding electrons further away from the nucleous Polarity of bond from data booklet calculate ΔE

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Shapes of molecules (VSEPR)

VSEPR

Shapes of molecules Negative charge centers 2

: Linear

3

:Trigonal planar

4

:Tetrahedral

5

:Trigonal bipyramid (- equatorial)

6

:Octahedral – square bipyramid (axial)

find number of electron pairs/charge centres in valence shell of central atom;

electron pairs/charge centres in valence shell of central atom repel each other;

SO2:bent 117O from 120O H2O: bent 104.5O from 109.5O Diamond, graphite and C60 fullerene

to positions of minimum energy/repulsion / maximum stability;

Bonding Graphite and C60 fullerene: covalent bonds and van der Waals’/London/dispersion forces; Diamond: covalent bonds (and van der Waals’/London/dispersion forces);

pairs forming a double or triple bond act as a single bond;

non-bonding pairs repel more than bonding pairs

Delocalized electrons Graphite and C60 fullerene: delocalized electrons; Diamond: no delocalized electrons; Structure Diamond: network/giant structure / macromolecular / three-dimensional structure Graphite: layered structure / two-dimensional structure / planar; C60 fullerene: consists of molecules / spheres made of atoms arranged in hexagons/pentagons; Bond angles Graphite: 120o Diamond: 109o; C60 fullerene: bond angles between 109–120o;

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Hybridization Graphite: sp2

Diamond: sp3

Fullerene: sp2 and sp3.

Number of atoms each carbon is bonded to Graphite and C60 fullerene: each C atom attached to 3 others; Diamond: each C atom attached to 4 atoms / tetrahedral arrangement of C (atoms) Structure and bonding in Si macromolecular tetrahedral structure Structure and bonding in SiO2 macromolecular; each Si bonded covalently to 4 oxygen atoms and each O atom bonded covalently to 2 Si atoms / single covalent bonds; have tetrahedral structure.

Structure and bonding in XeO2 macromolecular; each Xe bonded covalently to 2 oxygen atoms and each O atom bonded covalently to 2 Xe atoms / single covalent bonds;

xenon dioxide has two non-bonding pairs of electrons; xenon-oxygen bonds will have a square planar distribution Intermolecular forces:  London Temporary dipole forces due to momentary unevenness in spread of electrons. Weakest of intermolecular forces. Increase with increasing molar mass.  Dipole – Dipole Permanent electrostatic forces of attraction between polar molecules. Stronger than van der Waals’  Hydrogen bonding Occurs when hydrogen attached to a highly electronegative element (N, F, or O) is bonded to another highly electronegative element (N, F, or O). Stronger than dipole:dipole forces.

Physical properties Melting and boiling points:

Solubility :

Conductivity:

Covalent macromolecular structure Intermolecular forces

Intermolecular forces Like dissolve like

Metallic bonding

Ionic bond

Graphite/fullerenes Molten / aqueous solutions of ionic compounds

Metallic bonding Group 1 vs group 7 IB Chemistry Revision –Dr. D. Bampilis

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Sigma (σ) and pi (π) bonds. sigma (σ) axial overlap with electron density between the two carbon atoms/nuclei pi (π) parallel overlap of p orbitals with electron density above and below internuclear axis so shared electrons are above and below internuclear axis/ σ bond

σ bonds vs π bonds -axial

-

wide side

-first

-

second

-stronger -

weaker

Hybridization. mixing of (atomic) orbitals to form molecular orbitals of equal energy lower in total energy than atomic orbitals Resonance: when several Lewis structures are used collectively to describe the actual molecular structure. The actual structure is an approximate intermediate between the canonical forms, but its overall energy is lower than each of the contributors Contributing structures differ only in the position of electrons, not in the position of nuclei.

Delocalization of electrons A phenomenon in which valence electrons provided by individual atoms are no longer held in the near vicinity of that atom, but are mobile and shared by a number of atoms. Occurs (1) in metals(and graphite), where electrons can move throughout the entire crystal structure; (2) in organic compounds that have alternate double and single carbon-carbon bonds, orientated in such a way that p-electron can overlap, providing a pathway for electron movement; (3) in certain inorganic species, such as nitrate and carbonate ions, where p-orbital overlap can occur.

Delocalization stabilizes a structure, giving it a lower enthalpy than it would have if double and single bonds are arranged in such a way that orbital overlap cannot occur.

IB Chemistry Revision –Dr. D. Bampilis

BeX2, BeH2: sp hybridization BX3, BH3: sp2 hybridization O=C=O ,C=C=C (sp hybridization): linear

Different bond length C-O: CH3COOH vs CH3COODifferent bond length N-O: HNO3 vs NO3-, HNO2 vs NO2-

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Periodic table Blocks Group: (elements in vertical) columns in periodic table elements in same group have similar chemical properties group number gives number of valence/outer shell electrons for main group elements Period: (elements in horizontal) rows in periodic table elements within a period, atoms have same number of shells/energy levels (but number of electrons in valence/outer shell increases). period number gives same number of shells/energy level Atomic radius: Half the distance between the nuclei of two bonded atoms of the same element.

The trend in atomic radius across group increases down to group â&#x20AC;&#x201C; more shells The trend in atomic radius across period 3 decreases (from left to right/across period 3); same number of shells/energy levels / shielding remains the same; number of protons/nuclear charge increases so attraction of nucleus on outer electrons increases Ionic radius is the radius of an atom's ion.

Compare the ionic radius: anions in the same period with electron configuration of the next noble gas: increases to the left cations in the same period with electron configuration of the previous noble gas: increases to the left ions with the same electron configuration of a noble gas: increases with atomic number decreases

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Electronegativity the ability of an atom to attract a bonding pair of electrons; increases from right to left across a period.

Non-metals generally have higher electronegativity values than

metals.

decreases on descending a group. halogens have the highest electronegativity values.

Noble gases are not assigned electronegativity values do not react / do not attract electrons / stable electron configuration / full outer electron shell / do not form bonds; Why electronegativity increases across period 3 in the periodic table increasing number of protons / atomic number / nuclear charge; atomic radius decreases same shell / similar shielding from inner electrons; Trends in properties  SMALLER ATOMIC RADIUS SMALLER IONIC RADIUS(cations or anions) HIGHER ELECTRONEGATIVITY HIGHER IONIZATION ENERGY LOWER METALLIC CHARACTER HIGHER MELTING POINT OF THE METALS

LARGER ATOMIC RADIUS LARGER IONIC RADIUS LOWER ELECTRONEGATIVITY LOWER IONIZATION ENERGY HIGHER METALLIC CHARACTER LOWER MELTING POINT OF THE METALS HIGHER MELTING POINT OF THE HALOGENS

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Chemical properties of alkali metals · The reactivity increases down the group · Reaction with water M(s) + 2H2O(l)

2M+(aq) + 2OH-(aq) + H2(g)

alkali metals are dangerous since it is flammable when it forms hydrogen on contact with water · Reaction with halogens 2Na(s) + Cl2(g)

2NaCl(s) (redox reaction)

·Reaction with O2 4Li(s) + O2(g)

2Li2O(s) normal oxide

· Good reducing agents Chemical properties of halogens · Good oxidizing agents · The oxidizing ability decreases down the group (F2>Cl2>Br2>I2) Cl2(aq) + 2Br-(aq)

2Cl-(aq) + Br2(aq)

Cl2(aq) + 2I-(aq)

2Cl-(aq) + I2(aq)

Br2(aq) + 2I-(aq)

2Br-(aq) + I2(aq)

· Test for halide ions Ag+(aq) + X-(aq)

AgX(s):

AgCl white, AgBr cream, AgI yellow The trend in reactivity of the halogens. reactivity decreases down group as atomic radius increases / more electron shells; attraction of nucleus on electrons decreases / electron affinity decreases; Chemical properties The oxides of period 3 elements From metal to non – metal From ionic (Na2O, MgO, Al2O3) to covalent (SiO2, P4O10/P4O6, SO3/SO2, Cl2O7/Cl2O)

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From basic to acidic of the oxides Basic: 2Na+(aq) + 2OH-(aq)

Na2O(s) + H2O(l) Na2O(s) + 2HCl(aq)

2NaCl(aq) + H2O(l)

MgO(s) + H2O(l)

Mg(OH)2(aq)/(s)

MgO(s) + 2HCl(aq)

MgCl2(aq)+H2O(l)

Amphoteric: Al2O3 is insoluble in water Al2O3(s) + 6HCl(aq)

3AlCl3(aq) + 3H2O(l)

Al2O3(s) + 3H2SO4(aq)

Al2(SO4)3(aq) + 3H2O(l)

Al2O3(s) + 2NaOH(aq)

2NaAlO2(aq) + H2O(l)

or Al2O3(s) + 2NaOH(aq) + 3H2O(l)

2NaAl(OH)4(aq)

Acidic: SiO2(s) + 2NaOH(aq)

Na2SiO3(aq) + H2O(l)

Al2O3 and SiO2: do not react with water P4O6(s) + 6H2O(l)

4H3PO3(aq)

P4O10(s) + 6H2O(l)

4H3PO4(aq)

SO2(g) + H2O(l)

H2SO3(aq)

SO3(g) + H2O(l)

H2SO4(aq)

Cl2O(g) + H2O(l)

2HClO(aq)

Cl2O7(g) + H2O(l)

2HClO4(aq)

The clorides of period 3 elements From ionic (NaCl, MgCl2) to covalent (Al2Cl6, SiCl4, PCl3/PCl5, S2Cl2, Cl2) The reaction with water: NaCl(s) MgCl2(s)

H2 O

Na

6H2 O( l)

(aq)

Cl

(aq)

Mg H2 O

dissolve in water (neutral solution) 2

6

(aq)

2Cl

IB Chemistry Revision â&#x20AC;&#x201C;Dr. D. Bampilis

(aq)

dissolve in water Page 16

(acidic solution: [Mg(H2O)6]2+ [Mg(H2O)5(OH))]++H+)

Al2O3 (ionic) Al2Cl6 (covalent)

2AlCl3(s) + 3H2O(l)

Al2O3(s) + 6HCl(aq) Al2O3 amphoteric

SiCl4(l) + 4H2O(l) SiCl4(l) + 4H2O(l)

Si(OH)4(aq) + 4HCl(aq) or SiO2·2H2O(s) + 4HCl(aq)

PCl3(l) + 3H2O(l)

H3PO3(aq) + 3HCl(aq)

PCl5(l) + 4H2O(l)

H3PO4(aq) + 5HCl(aq)

Cl 2(aq)

H 2 O (l)

HOCl(aq)

AlCl3 acidic

HCl(aq)

Transition metal: An element the atoms of which have an incomplete set of d-electrons in their penultimate shell in one or more of their oxidation states. This includes all elements in the

Common redox number of transitions metals: +2

d-block except those with complete d-orbitals, such as zinc. Scandium is sometimes excluded from the transition metals because its ions have empty d-orbitals, and thus do not exhibit transition characteristics.

Characteristic properties of transition elements. partially filled d subshell variable oxidation numbers form complex (ions) structure of the H2O2

form coloured compounds

H-O-O-H

catalytic (behaviour) CH 2 3H 2(g)

CH 2(g) H 2(g) N 2(g)

2H 2 O 2(aq)

Fe(s)

MnO 2( s )

2SO2(g) O2(g)

Ni

CH3CH3(g)

2NH 3(g) (Haber process) 2H 2 O (l)

V2 O5(s)

Manganese(IV) oxide is blackish or brown solid

O 2(g)

2SO3(g) (Contact process)

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Ligand species with lone/non-bonding pair (of electrons); which forms dative covalent bond to metal ion (in complex); Lewis base which forms coordinate/dative covalent bonds with a metal ion (which acts as a Lewis acid); six coordinate complexes: are octahedral 3+

3-

Ligand H2O, F-, Cl-, Br-, I-, NH3, CN-, PCl5

Transition metal

,[Fe(H2O)6] , [Fe(CN)6]

complexes do not

four coordinate complexes: are tetrahedral,

obey VSEPR

[CuCl4]2- or square planar, [Ni(CN)4]2two coordinate complexes: are linear [Ag(NH3)2]+ The complex of transition elements is coloured d sub-level partially occupied;

Sc â&#x20AC;&#x201C; Zn are not

d orbitals split (into two sets of different energies);

transition metals.

frequencies of (visible) light absorbed by electrons moving from lower to higher d levels;

Cu+: are colourless

Sc3+ - Zn2+ - Ti4+ -

colour due to remaining frequencies / complementary colour transmitted; changing the ligand / coordination number / geometry changes the amount the d orbitals are split/energy difference between the d orbitals

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Stoichiometry

Finding moles: n

Empirical formula from percentage composition

n

Determining empirical formulas · from the mass of each element ( oxygen indirectly mO=mORG.COM.-mC-mH) · from the percentage composition The E.F. of alkenes: CH2, the

n n n

m (m : gr) M c V (solutions, V : dm 3 ) N L V (gases, V : dm 3 ) Vm PV (gases) RT

M.F.;(CH2)n Molecular formula from Empirical formula and Mr Mr: no units – Molar mas: g/mol Ideal gas (T↑, P↓)  n

V Vm

Ideal gas equation T in K

 PV

nRT, PV

m RT, M r Mr

mRT , Mr PV

dRT (gas) P

P: Pa V: m3 or P: kPa V: dm3 T:K

n: mol

R: 8.31 J·K-1mol-1

Solutions Concentration: n or n = cV c: V concentration of solution (mol dm-3) n: number of mole of solute (mol) V: volume of solution (dm3) Total solution – sample solution: same concentration Dilution of solutions:c1V1 = c2V2 Mix of 2 solution with same solute: c1V1 + c2V2 =cF(V1+V2) c=

IB Chemistry Revision –Dr. D. Bampilis

You must be clear which type of particle you are considering.

Do you have one mole of atoms, molecules or ions?

Do not confuse the mass of a single molecule with the mass of one mole of a substance.

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1 mol AxBy x mol A

y mol B

Stoichiometry: mole ratio – Volume ratio (for gases) -Work out the number of moles of anything you can -Use the chemical equation to work out the number of moles of the quantity you require.

-Convert moles to the required quantity.

2NaOH(aq) + H2SO4(aq) Na2SO4(aq) + 2H2O(l) : nNaOH=2nH2SO4 Limiting reactant – reactant in excess Percentage yield exp erimantal yield 100 percentage yield theoritical yield 1 Stoichiometric ratio – mole ratio : mol of reactants and products not of initial and final mole (chemical equilibrium) Mixture (x mol, y mol) mmix = xM1 + yM2, gases: n mix

V Vm

Water of crystallization Back titration (excess) Linked reactions

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Uncertainties

Accuracy % difference

experimental value accepted value 100 accepted value 1

Precision Closeness of agreement of a set of measurements to each other / Allow reproducibility/consistency of measurement / measurements with small random errors/total amount of random errors/standard deviation / a more precise value contains more significant figures Mistakes can be avoided â&#x20AC;&#x201C; errors can be minimized Systematic error equipment

Poor accuracy : change

poorly calibrated instruments badly made instruments instrument parallax error Error in reading an instrument employing a scale and pointer because the observer's eye and pointer are not in a line perpendicular to the plane of the scale.

Systematic error: liquid remaining in measuring cylinders not all solid transferred precision uncertainty of stopwatch ability of human eye to detect colour change

poorly timed actions Random error:

Random uncertainty experiments

Poor precision: more

mass uncertainties

synchronizing mixing and starting timing

volume uncertainties

uncertainty of concentrations of solutions

time uncertainties

temperature of solutions

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Significant figures leading zeros 0.023=2.310-2: 2 significant figures captive zeros 5.008: 4 significant figures trailing zeros 0.300: 4 significant figures 4102: 1 significant figures 400: 4.00102: 3 significant figures Multiplication and division 2.54 2.6 6.604 6.6

Addition and subtraction 3.647 4.5 3.66 11.627 11.6

Absolute uncertainty Percentage uncerta int y

adsolute uncerta int y 100 measurement 1

Addition: Add absolute uncertainties Multiplication, power: Add percentage uncertainties Graphical techniques directly propotional (y=mx+c) inversely propotional

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Energetics Kinetic molecular theory Molecules have no volume Gas molecules exert no force on each other unless they collide Collisions of molecules with each other or the walls of the container do not decrease the energy of the system The molecules of a gas are in constant and random motion The temperature of a gas depends on its average kinetic energy Κ.Ε. Τ Exothemric reaction system produces heat /standard enthalpy change negative , ΔΗ=HP – HR < O / bond formed more energetic than bond breaking / more stable product Endothermic reaction system absorbs heat from surroundings / standard enthalpy change positive /bond breaking more energetic than bond formation ΔΗrxn: state are very important (g,l,s) - ΔHf H2O(g) vs H2O(l) Calorimetry A device for measuring enthalpy changes for reactions. In a simple calorimeter all the heat evolved in an exothermic reaction is used to raise the temperature of a known mass of water.

Q= m c ΔΤ m: the mass of water (liquid) – not reactant or product ΔQsol = heat absorbed by solution + heat absorbed by calorimeter = (msolution  csol  ΔΤsol) + (mflask  cflask  ΔΤflask)=(mc+C)ΔT Assumptions: all heat is transferred to water/ solution / no heat loss; specific heat capacity of reactants and products are zero/negligible / no heat is absorbed by reactants and products; density of water/solution =1.0/ density of solution = density of water; mass of H2O= Volume of H2O heat capacity of cup is zero / no heat is absorbed by the cup; specific heat capacity of solution = specific heat capacity of water; temperature uniform throughout solution; incomplete combustion water forms as H2O(l) instead of H2O(g) ; IB Chemistry Revision –Dr. D. Bampilis

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Q = n ΔH

Stoichiometry: Q = n ΔH To compensate the heat lost:

Q= m c ΔΤ

insulate the reaction vessel - use a lid -draw a temperature versus time graph

ΔΤ=Δθ

ΔΤ from extension of the curve at the time – start the reaction Hess’ law Enthalpy change for a reaction depends only on difference between enthalpy of products and enthalpy of reactants. It is independent of pathway.

Enthalpy of Combustion: The energy released when one mole of a compound is burned in excess oxygen.

Average bond enthalpy energy required to break 1 mol of a bond in a gaseous state; energy released when (1 mol of) a bond is formed in a gaseous state / enthalpy change when (1 mol of) bonds are formed or broken in the gaseous state.

ΔΗrxn = ΣΔΗc(reactants) – ΣΔΗc(products) ΔΗc<0 (1 mol) ΔΗc0 organic compound H2O(l)

average values obtained from a number of similar compounds Average bond enthalpy (different from real / gas state- average- resonance structures)

ΔΗrxn= ΣD(bond broken) ΣD(bond formed) (gas state)

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The reaction HCl(g)

1 2

H2(g) +

1 Cl2(g) 2

does

not represent the bond enthalpy of HCl Standard enthalpy change of formation, ∆Hf heat/enthalpy change/required/absorbed when 1 mol of a compound is formed from its elements in their standard states/at 100 kPa/298 Κ ΔΗf (1 mol) ΔΗf element=0

ΔΗrxn = ΣΔΗf(products) – ΣΔΗf(reactants) ∆Hf of CO2 =∆Hc of C ∆Hf of H2O =∆Hc of H2

Standard enthalpy change of neutralization ΔΗn<0 Lattice enthalpy Energy required to change 1 mole of a solid ionic compound crystal into its gaseous ions or vice versa U

k

q1 q 2 r

Electron affinity The enthalpy change when 1 mol of atoms in gaseous state gains 1 mol of electrons.

Born – Haber cycle a. Atomization of the solid metal, ΔΗoat > 0 b. Ionization of the gaseous metal, ΔΗοI >0 c. Dissociation or atomization of the molecular non – metal into atoms, ΔΗoD > 0

(EA<0 ΔΗlfo<0)

d. Addition of electrons to the non – metal atoms (electron affinity) ΔΗοΕΑ1 < 0, ΔΗοΕΑ2> 0. e. Formation from the gaseous ions the solid ionic lattice, ΔΗolt (opposite lattice enthalpy) ΔΗolf < 0 More negative standard enthalpy, change of formation of the ionic compound, the more stable the compound will be.

Lattice enthalpy (ch arg e of cation) (ch arg e of anion) k sum of ionic radii

IB Chemistry Revision –Dr. D. Bampilis

The effect of charge is larger than the effect of the size

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Lattice enthalpy is exothermic as the making of the lattice is endothermic as the breaking of the lattice Difference between experimental value of ΔΗοLf and theoritical value covalent character Entropy A measure of the disorder of a system

ΔSrxn = ΣS(products) – ΣS(reactants)

Predicting the entropy change

(J

kJ)

ΔS>0: increase the disorder · mixing different types of particles · melting – evaporation – sublimation · increasing temperature · increasing the number of particles Predicting the entropy change - states of reactants and products - relative number of mole of reactants and products ΔGrxn =

- gaseous reactants/products

ΣnΔG(products) – Gibbs free energy

ΣmΔG(reactants) 0

0

o

S.T.P for ΔΗ / ΔG 1 atm, 25 C, 1 M for Vm 1 atm=101kPa, 273 K or 00 C o

Spontaneous ΔG <0

ΔG = ΔH – ΤΔS Spontaneous ΔGo<O

ΔSorxn

ΔHorxn

ΔGorxn

spontaneous

+

-

-

Yes

+

+

+ or -

yes, high temperatures

-

-

+ or -

yes, low temperatures

-

+

+

no

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Kinetics Rate of reaction change in concentration of reactant/product with time / rate of change of concentration; (initial, average, instantaneous, experimental methods for measuring) Rate N2(g) + 3H2(g)

2NH3(g)

3RN2=RH2

Activation energy, Ea minimum energy needed (by reactants/colliding particles) to start a reaction / for a successful collision; energy difference between reactants and transition state Maxwell – Boltzmann curve Graph showing the distribution of kinetic energies among molecules.

Collision theory · Collide

·E

Ea

· Correct geometry

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Explain the increase in rate in terms of collision theory and Maxwell-Boltzmann energy distribution curves kinetic energy of molecules increases; frequency of collision increases; greater proportion of molecules have energy greater than/equal to activation energy; correct Boltzmann-energy distribution curves showing curve at higher temperature on the right side; broadening of the curve; Factors that can affect the rate of a reaction The state of the reactants The activation energy The frequency of the collisions between particles The average kinetic energy of the particles

catalyst concentration/ pressure; particle size / surface area; temperature light; temperature

The effect of temperature more collisions Ea – Maxwell – Boltzmann: larger propotion of particles have energies in excess of the activation energy for the reaction (more important) 10oC – double The effect of concentration As the concentration increases, the reaction rate increases. The number of collision that occur per unit time has increased.

The effect of surface area Increasing surface area increases the area for collisions.

The effect of the catalyst on the reaction rate increases rate of forward and reverse reactions (equally) provides a reaction pathway with lower activation energy; reactants adsorb onto the catalyst surface and bonds weaken resulting in a decrease in activation energy more molecules/particles have sufficient energy to react; IB Chemistry Revision –Dr. D. Bampilis

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Heterogeneous catalysts gases adsorb on surface of catalyst / provides surface for the reaction; lowers activation energy / provides alternative pathway for reaction; Order of reaction – Rate Law –Reaction mechanism Rate expression: Rate = k[R1]m[R2]n (experimental ) Mechanism: rate determining step-slowest-k↓-Ea↑ The units of k are:(mol.dm-3)(1-OVERAL ORDER).time-1 Reaction mechanism Must agree with the overall stoichiometric equation A maximum of two particles can react in any one step All species in the rate equation must appear in the mechanism in or before the rate determining step.

The power of a particular reactant’s concentration in the rate equation indicates the number of times it appears in the mechanism up to and including the rate determining step.

Termolecular reactions (more complicated reaction) steps: i.

Fast chemical equilibrium

ii.

Slow reaction

Molecularity is the number of particles in the slowest / rate determining step of the reaction; Molecularity (SN1-SN2): can only be applied to elementary reactions.

Note: molecularity and order are not the same Activation energy – transition state : (activated complex) – Intermediates Catalyst: can be appear in the rate equation -maybe is involved in the rate determining step, but is regenerated in the second step - do not appear in the overall chemical equation Transition state (activated complex) : the highest energy species on the reaction pathway

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Intermediate: produced in step 1 and used up in step 2 , do not appear in the rate equation.

Graphs rate / concentration and concentration / time Graphs showing reaction orders The results of an experiment investigating the effect of the concentration of a reactant, we draw a diagram of rate or 1/time vs concentration of a reactant The results of an experiment investigating the change of the concentration of a reactant with time, we draw a diagram of concentration of a reactant vs time Half-life the time taken for the concentration/amount (of ) to decrease to half its original value

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Zero order reaction Rate=k[A]o

Rate is constant.

The concentration time graph of reactant is a straight line with a negative gradient First order reaction The rate â&#x20AC;&#x201C; concentration graph is a straight line with gradient equal to k

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For a first order reaction: d[A] dt

k[A]

At t1/ 2 :[A]

d[A] [A]

kdt

int egration

ln[A]o

ln[A]

kt

1 [A]o 2

The expression becomes

kt1/ 2

ln[A]o

1 ln [A]o 2

ln 2

t1/ 2

0, 693 k

Second order reaction Rate=k[A]2 The rate – concentration graph is a curve – parabolic, because it depends on the square of concentration.

The concentration – time graph is a curve, the rate (the gradient) decreases much more quickly than the first order reaction.

t1/ 2

0, 693 for first order K (x1/ 2) for zero order

t1/ 2

(x2) for sec ond order

t1/ 2

Half-time

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A plot of log rate against log[A] gives a straight line graph whose intercept is the value for log k and the gradient is equal to the order of the reaction.

Arrhenious equation k

Ae

Ea / RT

or ln k

ln A

Ea RT

A: contains information related to the frequency and the orientation of the collisions : represents the fraction of collisions that have E>Ea The rate constant will increase as the absolute temperature increase(not linearly).

The rate constant will increase as the Ea decrease.

If a catalyst is added gradient of the line less steep (less negative).

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Equilibrium

Chemical equilibrium reversible reaction rate of forward reaction equals rate of backward reaction; concentrations of reactants and products do not change / constant macroscopic properties; Closed system Equilibrium Law

Le Chatelier’s Principle If an equilibrium is subjected to a stress, the equilibrium will shift to minimize the effect of the stress.

Temperature: is the only factor that can affect both the equilibrium position and the equilibrium constant.

Exothermic

Endothermic

Temperature change

Kc

Kc

Increase

Decreases

Increases

Decrease

Increases

Decreases

If ΔHRXN=0, change in temperature will have no effect whatsover on the position of the equilibrium.

The heat change in the formation of an ester is almost zero, and changes to temperature have very little effect on the position of equilibrium.

Concentration (change of mole) Increasing the concentration of a reactant, the position of equilibrium moves to right.

Increasing the concentration of a product, the position of equilibrium moves to left.

Decreasing the concentration of a reactant, the position of equilibrium moves to left.

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Decreasing the concentration of a product, the position of equilibrium moves to right.

CaCO3(s)

CaO(s) + CO2(g):adding or removing solid without any change

in the volume of the reaction, we have no change at the position of the equilibrium Changes to physical equilibrium systems Br2(l)

Kc

Br2(g)

[Br2(g) ]

H 2 O(s)

Kc

H 2 O(l)

[H2 O(l) ]

Vapour Equilibrium Vapour pressure: kp=P0 -depends only on the nature of the liquid and the temperature -T↑ P↑( exponentially-difference from ideal gases equation) -High vapour pressure, more volatile liquid, lower boiling point Pressure (Change of volume – for gases) An increase in external pressure the reaction moves toward the side with the lower number of particles in the gas state.

A decrease in external pressure the reaction moves toward the side with the greater number of particles in the gas state.

If the equilibrium system has the same number of reactant gas particles as the products gas particles, a change in pressure will have no affect whatsoever on the position of equilibrium.

Dilution of aqueous solutions If water is added, the concentration of all species will decrease. The position of equilibrium moves to the side of the reaction with the greater number of mole of particles.

The role of a catalyst The addition of a catalyst does not change the position of equilibrium or the value of the equilibrium constant. Equilibrium is simply reached more quickly.

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Two economic benefits of using the catalyst to speed up the reaction reduces energy costs (as reaction occurs at lower temperatures); catalyst can be reused; increases reaction rate so more product produced in a given time For the system N2 (g)+ 3H2(g) 2NH3 (g) state and explain the effect on the position of equilibrium of adding some helium gas but keeping the total gas volume constant.

No change / The total pressure would increase but the concentration of products and reactants remain the same thus the equilibrium position does not change.

Haber process: N2(g) +3H2(g)

2NH3(g)

200 atm (pressure) 700 K (temperature)

Haber process pressure is lower/moderate and temperature is higher in Haber process Pressure: high pressure shifts equilibrium to right; high pressure (faster rate but) expensive /dangerous Temperature: low temperature shifts equilibrium (even further) to right; low temperature gives slower rate (but high yield); high pressure increases yield and lower temperature decreases rate; Qc / kc Qc

Concentration fraction

P A

p in a in

Q B

q in b in

IB Chemistry Revision â&#x20AC;&#x201C;Dr. D. Bampilis

Contact process: 2SO2 + O2

2SO3

450OC â&#x20AC;&#x201C; 2 atm - 99%

If Qc<Kc : The reaction move to right.

If Qc=Kc : Equilibrium.

If Qc>Kc : The reaction move to left.

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Acids – Bases Bronsted – Lowry conjugate pair differ by one H+ Acid:proton/H+ donor; Base: proton/H+ acceptor Lewis dative bond Acid electron pair acceptor; Base electron pair donor; Weak acid and strong acid. weak acids partially dissociated/ionized and strong acids completely dissociated/ ionized (in solution/water) Strong acids: HCl, HBr, HI, H2SO4, HNO3 Strong bases: NaOH, KOH, Mg(OH)2 Experimental distinguish between strong and weak acid with the same concentration: – better conductivity – lower pH – more violent reaction with Me or carbonates – more heat neutralization realized Properties of acids and bases Reactions of acids  React with metals above hydrogen in the reactivity series (not Cu, Ag or Hg) to produce a salt and hydrogen gas e.g. 2HCl(aq) + Mg(s)

MgCl2(aq) + H2(g)

 React with metal hydroxides to produce a salt and water (neutralization) e.g. 2HCl(aq) + Mg(OH)2(s)

MgCl2(aq) + 2H2O(l)

 React with NH3 to form salt (neutralization).

NH3(aq) + HCl(aq)

NH4Cl(aq)

 React with metal oxides to produce a salt and water (neutralization) e.g. 2HCl(aq) + MgO(s)

MgCl2(aq) + H2O(l)

 React with metal carbonates to produce a salt, water and carbon dioxide e.g. 2HCl(aq) + MgCO3(s)

MgCl2(aq) + H2O(l) + CO2(g)

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 React with metal hydrogen carbonates to produce a salt, water and carbon dioxide e.g. 2HCl(aq) + Mg(HCO3)2(s)

MgCl2(aq) + 2H2O(l) + 2CO2(g)

Reactions of bases  Alkalis react with acids to produce a salt and water (neutralization) e.g. NaOH(aq) + HCl(aq)

NaCl(aq) + H2O(l)

 Metal oxides react with acids to produce a salt and water (neutralization) e.g. MgO(s) + 2HCl(aq)

MgCl2(aq) + H2O(l)

 Metal carbonates react with acids to produce a salt, water and carbon dioxide e.g. Na2CO3(s) + 2HCl(aq)

2NaCl(aq) + H2O(l) + CO2(g)

 Metal hydrogen carbonates react with acids to produce a salt, water and carbon dioxide e.g. NaHCO3(s) + HCl(aq)

NaCl(aq) + H2O(l) + CO2(g)

 Displacement of ammonia from ammonium salts NH4Cl(s) + NaOH(aq)

NaCl(aq) + H2O(l) + NH3(g)/(aq)

 Ammonia reacts with acids to produce an ammonium salt e.g. NH3(aq) + HCl(aq)

NH4Cl(aq)

Acid and base dissociation constants Relative strengths of acids and bases: Ka, Kb, pKa and pKb: Ka

Kw Kb

strong: Ka or Kb max strong: pKa or pKb min pH or [H3O +] can be used to compare acid strength only if equal concentrations of acids are being compared Self-ionization of water - The pH scale-kw ka, kb, kw : assumption that the concentration of water, [H2O], is constant Kesterification, Kphase change: [H2O] appeared T↑ kw↑(endothermic) pure water or neutral solution pH= pOH= pkw: 250C : [H3O+] =[ OH-]=10-7 – pH= pOH=7 θ 250C : [H3O+] =[ OH-] 10-7 – pH= pOH 7 IB Chemistry Revision –Dr. D. Bampilis

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θ 250C : [H3O+] =[ OH-] 10-7 – pH= pOH 7 Calculating pH of weak acid and base making an ICE (initial, change, equilibrium) table, you can find unknown concentration values that can be plugged into equilibrium expression Ka or Kb. Assumptions: temperature 298 K, the concentration of H3O+ or OH- from ionization of H2O is negligible.

Buffer solution. solution which resists change in pH when small amounts of acid or base or water are added; Explanation of buffer action (neutralization) acid 'removes' most of any added hydroxide ions: -

-

CH3COOH(aq) + OH (aq)

CH3COO (aq) + H2O(l)

or +

NH4

-

(aq)

+ OH (aq)

NH3(aq) + H2O(l)

base 'removes' most of any added hydrogen ions: -

CH3COO (aq) + H

+

CH3COOH(aq)

(aq)

or NH3(aq) + H

+

+

NH4

(aq)

(aq)

The pH of a buffer solution For the equation: HA (aq) H 2 O(l) pH

pK a

log10

salt acid

pK b log10

salt base

(aq)

A

(aq)

When [acid]=[salt]:[H3O+]=Ka and pH=pKa

For the equation B(aq) H 2 O(l) pOH

H3O

HB

(aq)

OH

(aq)

When [base]=[salt]:pOH=pKb

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Other buffer solution: NaHA – Na2A(exception NaHSO4 – Na2SO4) H2A – NaHA(exception H2SO4 – NaHSO4) CH3COONH4 Amino acids Comparing the pH of solution NH3with buffer solution NH3/ NH4Cl (Le Chatelier’s Principle) The buffer has lower pH NH3(aq) + H2O(l) ⇄ ΝΗ4+ (aq) + ΟΗ- (aq) [ΝΗ4+] ↑ shift the equilibrium to the left so the [ΟΗ-] ↓ Salt hydrolysis Salts formed from the four possible combinations of strong and weak acids and bases.

Strong acid + Strong base Salt will be neutral; pH=7.0 Weak acid + Strong base Salt will be basic; pH>7 Strong acid + Weak base Salt will be acidic; pH<7 For weak acid + weak base Salt will be have pH depending on the relative strength of the weak acid and the weak base ( comparing the Ka, Kb ) Hydrolysis of the cations - results in strong attraction of the high charge density cations to the lone pair of one of six water molecules surrounding the ion, a process by which the water molecule loses a hydrogen ion, i.e. leaving the solution acidic(Mg2+ , Al3+ , TMx+ - depends on charge, ionic radius) Indicators HIn (aq) + H2O(l)

In (aq) +H3O+(aq)

Colour A Colour B + in acid/presence of H equilibrium lies to left (so colour A); in alkali/base/presence of OH– equilibrium lies to right (so colour B); Color change

The equation, derived from the acidity constan, states that when pH equals the pKa value of the indicator, both species are present in a 1:1 ratio.

Usually, the color change is not instantaneous at the pKa value, but a pH range exists where a mixture of colors is present. colour changes when [HIn (aq)] ≈ [In− (aq)] ;

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Titration curves Titration curve for a strong acid with a strong base.

Equivalence point: the point at which the acid and base have reacted in exact stoichiometric amounts.

Titration curve for a weak acid with a strong base.

Buffer region Half neutralized: pH=pKa

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Titration curve for a strong acid with a weak base.

Titration curve for a weak acid with a weak base.

Appropriate indicator for a titration End point of a titration is the point at which the indicator changes colour

Acid rain H2SO4 /HNO3/ H2SO3 /HNO2 /H2CO3 respiratory problems corrosion problems decomposition of ozone layer photochemical smog acidification of lakes damage to plants

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Redox Oxidation reduction in terms of oxidation number Oxidation: increase in oxidation number; Reduction: decrease in oxidation number Oxidation reduction in terms of electron transfer loss of electrons-gain of electrons; Oxidizing agents.

Reducing agents.

substance reduced / removes It is oxidized by losing electrons /gives electrons from another substance/ electrons to another substance / causes causes some other substance to be some other substance to be reduced oxidized F2, O3, Cl2, Br2, O2, I2, S

X2

2X

O2

2

O

MnO2

....

Mn 2

H H 2O 2

H 2O2

SO2

.... ( X : F , Cl , Br , I )

....

....

S 0 ....

H 2 SO4 (conc)

SO2 ....

HNO3 (dilute)

NO ....

HNO3 (conc)

NO2 ....

KMnO4

Mn 2

H

K 2Cr2O7

H

....

2Cr 3 ....

MnO4− → MnO2 + OH−

Fe

3

Fe

4

Sn2

Sn

K, Ba, Ca, Na, Mg, Al, Mn, Zn, Cr, Fe, Co, Ni, Sn, Pb, H2, Bi, Cu, Hg, Ag, Pt, Au.

C

CO2 ....,

S

H 2 SO4 ...., P

SO2

H 2 SO4 ....,

CO

CO2 ....,

H 2O2

O2 ....

2 HX

X 2 .... ( X

H2S

H 2 SO3

2 NaX

H 2O ....

H 3 PO4 ....

Cl , Br , I )

S 0 ....

H 2 SO4 ....

Na2 SO3

NaSO4 ....

X 2 .... ( X

2 NaXO

X 2 .... ( X

Na2 S

S ....

Fe2

Fe3

2

4

Cl , Br , I ) Cl , Br , I )

2

Sn KClO3

H2

Sn

....

KCl ....

IB Chemistry Revision –Dr. D. Bampilis

2 NH 3

N 2 ....

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Titration with KMnO4 When the students recorded the burette readings, following the titration with KMnO4 (aq), the top of the meniscus was used and not the bottom - Potassium permanganate has a very dark/deep (purple) colour so cannot read bottom of meniscus In an acidic solution, permanganate(VII) is reduced to the colourless +2 oxidation state of the manganese(II) (Mn2+) ion.

8 H+ + MnO4− + 5 e− → Mn2+ + 4 H2O Sulfuric acid used to accidified the solution, two other strong acids such as nitric acid, HNO3 (aq) or hydrochloric acid, HCl(aq), couldn’t be used HNO3 is an oxidizing agent HCl reacts with MnO4 – to form Cl2 In a neutral medium however, it gets reduced to the brown +4 oxidation state of manganese dioxide MnO2.

2 H2O + MnO4− + 3 e− → MnO2 + 4 OH−

Redox Equations Oxidation number: Cu:+2 , Charge: Cu:2+ Balancing redox half-equations: e-/H+/H2O or H2O /H+/ eHydrogen react as oxidizing agent when react with metals to form hydrides

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AN ACTIVITY SERIES FOR METALS Reactivity

Metals

Properties

K

Burn readily in oxygen to form oxides

Na

React with water to from hydrogen and hydroxides

Ca

increasing reactivity

Mg

React with acids to form hydrogen and salts

Al

Burn to form oxides if finely divided

Zn

React with steam to form hydrogen and oxides

Cr

React with cold acids to form hydrogen and salts

Fe Cd Co Ni Sn

Oxidize if heated in air or pure oxygen

Pb

No reaction with steam

Cu

Sn and Pb react with warm acids to form hydrogen and salts

Hg

Cu and Hg do not react with acids Ag

Do not react with oxygen

Pt

No reaction with steam

Au

No reaction with acids

Voltaic Cells Two different half-cells connected together to enable to electron transferred during the redox reaction to produce energy in the form of electricity. The electrons are produced at the half-cell that is most easily oxidized.

Half cell with more negative E0 will be the negative electrode in the cell and the electrons will flow from this half cell to the other one(oxidation â&#x20AC;&#x201C; anode).

Electrolytic Cells Used to make non-spontaneous redox reactions occur by providing energy in the form of electricity from an external source.

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Differences between an electrolytic cell and a voltaic cell Voltaic cells

Electrolytic cells

 Spontaneous exothermic reactions convert chemical energy into electrical energy.

 Non-spontaneous endothermic reactions require electrical energy, which is converted into chemical energy.

 Can be used as a power source.

 Requires an external power supply.

 Comprises two separate half-cells connected by a salt bridge and connecting wires.

 No need to separate half-cells and so no salt bridge is required.

 Electrons flow from negative electrode to positive electrode.

 Power supply forces electrons onto negative electrode and takes them from the positive electrode.

 The direction in which the positive ions flow in the salt bridge is opposite to the negative ions flow  Oxidation occurs at the anode.

 Oxidation occurs at the anode.

 Reduction occurs at the cathode.

 Reducing occurs at the cathode.

 Anode is negative.

 Anode is positive.

 Cathode is positive.

 Cathode is negative.

 Polarity of the electrodes is determined by the reactions occurring there.

 Polarity of the electrodes is determined by the power source.

 Reducing agent donates electrons to oxidizing agent via the external circuit.

 Power source supplies electrons to oxidizing agent and accepts electrons from reducing agent.

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Electrolysis The strongest oxidizing agent is reduced at cathode (negatively charged)

Electrolyte: cations move to negative electrode/cathode and anions to positive electrode/anode; Conductors: electrons flow from negative pole of battery to positive pole of battery Electrolysis of aqueous solutions: The products of an electrolysis reaction can be predicted from the electrochemical series.

We must consider the reactions for water/or H+ , OHCathode (-) : H 2 O (l)

1 H 2(g) 2

e

Electrolysis of NaCl: Concentrated

OH

(aq)

Eo

0.83V

(1mol dm-3):H2, Cl2 Dilute: H2,O2

Anode (+): 1 O 2(g) 2

The strongest reducing agent is oxidized at the anode (positively charged).

2H

(aq)

2e

H 2 O(l) E o

1.23V

Electrolysis with Reactive Electrodes In this type of cell, the electrodes are normal metals, not inert ones like platinum or carbon. ď&#x201A;&#x; The metal of the electrode needs to be considered at the positive electrode as a possible reducing agent.

There are 3 possibilities for Oxidation at the

Electroplating Anode (+):Me Solution: Me2(SO4)x or Me(NO3)x Cathode (-) : spoon

ANODE: 1. The anion in the solution is oxidized 2. Water is oxidized 3. The Anode Metal is oxidized

The cathode metal never reacts!!

Experimental conditions in electrolysis affect the rate increase in current/voltage/surface area of electrodes;

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Factors affecting the relative amounts of products formed during electrolysis Relative atomic mass of metal Current- time charge of ion (stoichiometry) m

Q=I·t

Ar Q zF

F = NA · qe- = 96.500 C

temperature of solution; surface area/size of electrode; material of electrodes; Electrolytic cells connected in series -the same current passes through each cell Standard electrode potential, EÖ the potential difference obtained when a half-cell is connected to a standard hydrogen electrode; under standard conditions / solute concentration of 1 mol dm–3 or 100 kPa for gases , 298 K measured relative to standard hydrogen electrode Explain the significance of the minus sign in E

Ö

The half equation is a reducing reaction The strongest reducing agent at the top right hand side The strongest oxidizing agent at the bottom left – hand side

the electrons flow from the half-cell to the standard hydrogen electrode / the half-cell forms the negative electrode when connected to the standard half-cell / Χ is a better reducing agent than H2 / Χ is above H2 in electrochemical series/ the half reaction is not spontaneous.

Calculating cell potentials If the value of Eº calculated is positive and greater than +0,3V (approximately) then the reaction is likely to occur.

If the value calculated is between 0 and +0,3V then its likely to be an equilibrium.

In the IBO Chemistry Data booklet the electrode potentials are arranged from negative to positive. The half equation higher up to the table is therefore the one that is reversed to give the overall spontaneous reaction.

Oxidizing agent with more positive Eo reacts with reducing agent with more negative Eo.

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E = E(reduction) E(oxidation) (in volts).

Spontaneous ΔGo < 0 Eo > 0

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Organic Chemistry Carbon â&#x20AC;&#x201C; Organic chemistry Features of a homologous series same functional group same general formula; difference between successive members is CH2; similar chemical properties; gradually changing physical properties Naming Naming esters: alkyl carbonate Founding E.T. M.T. CaCl2 absorb H2O - NaOH absorb CO2 Physical properties gradually changing in a homologous series Isomers compounds with the same molecular formula and different structural formula/different structures;

straight chain, increased surface area/more closely packed; stronger London forces

Exhaust gas

Reactions Combustion Hydrocarbon, CxHy, was burned in excess oxygen CxHy(g) + (x+y/4)O2(g) (y/2)H2O(g)

xCO2(g) +

Incomplete combustion Incomplete combustion occurs when the supply of air or oxygen is poor. Water is still produced, but carbon monoxide and carbon are produced instead of carbon dioxide.

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Local pollutant: CO volatile organics NO unburnt hydrocarbons NO2

Global pollutant: NO CO2 NO2

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Distinguishing between different organic compounds a)

test for alkenes

Alkenes decolorize bromine water because they undergo an electrophilic addition reaction with bromine: Add a few drops of bromine water to the sample and shake. If the bromine decolorizes, an alkene is present. If not, no alkene is present.

Distinguish between alkanes and alkenes alkanes â&#x20AC;&#x201C; no change / stays or turns brown; alkenes â&#x20AC;&#x201C; bromine (water) decolorizes; b)

test for haloalkanes

When haloalkanes are heated with dilute sodium hydroxide, a nucleophilic substitution reaction occurs and halide ions are produced. The halide ions can be identified using the tests described in AS Unit 2: Add aqueous sodium hydroxide to the sample and heat. Then allow to cool, add dilute nitric acid and then aqueous silver nitrate. A white precipitate soluble in dilute ammonia indicates that a chloroalkane was present, a cream precipitate soluble in concentrated ammonia indicates that a bromoalkane was present, and a yellow precipitate insoluble in ammonia indicates that an iodoalkane was present. c)

test for carboxylic acids

Carboxylic acids are acids and can liberate carbon dioxide from carbonates: Add sodium carbonate solution to the sample. If effervescence is seen, and the gas produced turns limewater milky, a carboxylic acid is present. d)

test for amines

Amines are basic.

Add universal indicator to the sample. If it turns blue/purple an amine is present. It will also have a fishy smell. e)

test for alcohols

Alcohols react with carboxylic acids in the presence of sulphuric acid to make esters.

Add ethanoic acid to the sample, followed by sulphuric acid and heat. If a the mixture starts smelling sweet and fruity an alcohol was present.

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Organic synthesis It is possible to make a large number of organic products from a few starting compounds and the necessary reagents.

1.

alkane  chloroalkane

reagents: chlorine conditions: UV light NB This reaction introduces a new functional group onto the molecule 2.

chloroalkane  alcohol

CH3CH2 CH3 + Cl2  CH3CH2 CH2Cl + HCl produce a mixture of products a poor yield of the desired product

R-Cl + NaOH  R-OH + NaCl

reagents: aqueous NaOH dilute conditions: warm, reflux 3.

chloroalkane  alkene

reagents: alcoholic KOH concentrated

CH3CH2Cl + KOH  CH2=CH2 + KCl + H2O

conditions: heat/reflux 4.

chloroalkane  nitrile

R-Cl + KCN  R-CN + KCl

reagents: aqueous KCN conditions: heat, reflux NB This reaction introduces an extra carbon atom onto the molecule 5. chloroalkane  primary amine

R-Cl + 2NH3  R-NH2 + NH4Cl

reagents: excess ammonia( to prevent production secondary amine see 6 reaction) conditions: heat

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6. primary amine  secondary amine

R1-NH2 + R2-Cl  R1R2NH + HCl

reagents: chloroalkane conditions: warm 7. secondary amine  tertiary amine

R1R2NH + R3-Cl  R1R2R3N + HCl

reagents: chloroalkane conditions: warm 8.

alkene  alkane

reagents: hydrogen, Ni catalyst conditions: 150 oC, 2 atm 9.

alkene  dibromoalkane

CH2=CH2 + H2  CH3CH3 Some margarine is made by hydrogenating carbon-carbon double bonds

CH2=CH2 + Br2  CH2BrCH2Br

reagents: bromine conditions: room temperature NB This reaction introduces a new functional group onto the molecule 10.

alkene  bromoalkane

CH2=CH2 + HBr  CH3CH2Br

reagents: hydrogen bromide conditions: room temperature 11.

alkene  alcohol

CH2=CH2 +H2O  CH3CH2OH

reagents: H2O conditions: concentrated sulfuric acid or phosphoric acid as catalyst 12.

nitrile  primary amine

R-CN + 2H2  R-CH2NH2

reagents: H2 , palladium / nickel catalyst. conditions: raised temperature and pressure

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13.

alcohol  alkene

Dehydration removal of H and OH from neighbouring carbon atoms;

CH3CH2OH  CH2=CH2 + H2O

reagents: concentrated sulfuric acid / concentrated phosphoric acid / hot Αl2Ο3/ hot ceramic conditions: heat, reflux 14. primary or secondary alcohol  carbonyl

R1R2CHOH + [O]  R1-COR2 + H2O

reagents: potassium dichromate and dilute sulphuric acid conditions: warm, distillation

Oxidation RCH2OH (distillation) 15. acid

RCHO

primary alcohol  carboxylic

R-CH2OH + 2[O]  R-COOH + H2O

reagents: potassium dichromate and dilute sulphuric acid conditions: heat, reflux

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16.

aldehyde  carboxylic acid

reagents: potassium dichromate and dilute sulphuric acid

R-CHO + [O]  R-COOH

conditions: heat, reflux 17. carboxylic acid  carboxylate salt

R-COOH + NaOH  R-COONa + H2O

reagents: NaOH conditions: room temperature 18. carboxylic acid + alcohol  ester reagents: concentrated sulphuric acid conditions: heat, organic solvent 19. acid anhydride + alcohol  ester conditions: warm 20. ester  carboxylic acid + alcohol reagents: concentrated sulphuric acid

R1COOH + R2OH R1COOR2 + H2O Use of esters: food flavouring/ perfumes / solvents / plasticizers/ glue (R1CO)2O + R2OH  R1COOR2 + R1COOH

R1COOR2 + H2O  R1COOH + R2OH

conditions: heat under reflux 21.

acid + primary amine 

N-substituted amide

R1-COCl + 2R2-NH2  R1-CONHR2 + R2-NH3Cl

Initially a salt is formed

conditions: heat at 2000C 22. carboxylic acid  acid anhydride

2R-COOH  (RCO)2O + H2O

conditions: heat

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Free radical substitution of alkanes Free radical substitution of alkanes: CH 4(g)

uv

Cl2(g)

CH3 Cl(g)

HCl(g)

The reaction mechanism: free radical substitution Initiation – increase in the number of free radical :

Cl Cl(g)

uv

2Cl (g)

Homolytic fission: the breaking of a covalent bond in which each atom involved in the bond retains one electron from the bond.

Propagation – no change in the number of free radical: Cl (g)

CH 4(g)

CH3 (g)

CH3 (g)

Cl2(g)

HCl(g)

CH3 Cl(g)

Cl (g)

Termination – decrease in the number of free radical: Cl (g)

Cl (g)

Cl (g)

CH3 (g)

CH3 (g)

Cl2(g)

CH3Cl(g)

CH3 (g)

CH3 CH3(g)

If there is sufficient chlorine then further substitution reactions will occur: CH3Cl(g)

Cl2(g)

uv

CH2 Cl2(g) uv

CH2 Cl2 (g) Cl2 (g) CHCl3 (l) Cl2 (g)

uv

HCl(g)

CHCl3 (l) HCl CCl4 (l) HCl

Nucleophilic substitution replacement of atom / group (in a molecule);by a species with a lone pair of electrons / species attracted to an electron-deficient carbon atom; The hydroxide ion is a better nucleophile than water OH– has a negative charge/higher electron density; greater attraction to the carbon atom (with the partial positive charge) Nucleophilic substitution reactions : SN1 faster than SN2 SN2 : bimolecular- transition state- primary RX nucleophile: :OH–, :CN–, :NH3 , the curly arrow must come from atom with lone pair of electon. (Anion↑ - Base↑ - Electronegativity↓) IB Chemistry Revision –Dr. D. Bampilis

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leaving group (Strength bond↓) substrate (hindrance) SN1: unimolecular- rate depends on [RX] only- intermediate carbocation – slow and fast step –tertiary RX nucleophile – no effect leaving group (strength bond↓) substrate (positive inductive effect) FACTORS AFFECTING SN1 AND SN2 REACTIONS OF HALOGENOALKANES Factor Effect on SN2 reaction rate Effect on SN1 reaction rate Nucleophile  Charged species are stronger No effect. The nucleophile does not nucleophiles than neutral species e.g. OH influence step 1, the rate-determining is a stronger nucleophile than H2O. heterolytic fission step.  Stronger bases are stronger nucleophiles, e.g. NH3 is a stronger nucleophile than H2O.

Halogen leaving group

 The less electronegative atom is a better nucleophile e.g. CN- is a stronger nucleophile than OH-.

The weaker the C-X bond the stronger the leaving group, e.g. the weaker C-I bond is more reactive than the stronger C-CI bond.

As for SN2. The C-X bond strength influences the ease of step 1, the rate-determining heterolytic fission step.

Good leaving groups are able to stabilize the charge on the transition state, the Ea is lower, means more rapid reactions.

Substrate

So RF, ROH, RNH2 do not normally undergo SN2 reactions.

The order of reactivity for halodenoalkanes is C-I > C-Br > C-Cl Primary halogenoalkanes show the least  Positive inductive effects by alkyl steric hindrance and so react faster than groups attached to the C-X carbon tertiary halogenoalkanes. stabilize the carbocation and so influence ease of reaction. Tertiary halogenoalkanes show the greatest inductive effect and so react faster than primary halogenoalkanes.

Rate: tert RX>secRX>primRX

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Halogenoalkanes react with : dilute NaOH(aq) to form ROH concentrated NaOH/Ethanol/reflux to form CnH2n Different isotopes will undergo hydrolysis at different rates? same rate as have same chemical properties different rate as molecules having different speeds/collision rate

Elimination H and X are eliminated from adjacent carbon atoms The mechanism first involves removal of a proton (abstraction) by the base.

This happens to the hydrogen atom which is on the carbon adjacent to the carbon holding the bromine atom. This hydrogen atom is said to be an 'alphahydrogen': OH- + CH3CHBrCH3

-

CH2CHBrCH3 + H2O

This produces a negatively charged species (a carbanion), which then rearranges by moving the negative charge lone pair into a pi orbital between the two carbon atoms with the loss of the bromine atom as a bromide ion: -

CH2CH2BrCH3

CH2=CHCH3 + Br-

E2 (bimolecular elimination)

 stronger base  high temperature It is possible that actual base is CH3-CH2O-, this ion is a stronger base than OHOH

CH3CH 2 OH

H 2 O CH3CH 2 O

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E1 (unimolecular elimination)

Basic nucleophiles favor the E1 reaction over the SN1 reaction Addition Polymerization (same composition)

The polymerization in the manufacture of plastics Polyethene household containers, plastic bags, water tanks and piping Polypropene clothing and thermal wear for outdoor activities Polychloroethene construction materials, packaging, electrical cable sheeting, carpets / gutters / rope / bottles it is one of the world's most important plastics economic importance plastics are cheap /a large industry / plastics have many uses / disposal of plastics plastics are not biodegradeable / plastics take up large amounts of space in landfill / pollution caused by burning of plastics Condensation polymerization reaction The economic importance of the condensation polymer production of nylon/clothes/carpets/ropes/Kevlar; Polyester dipole-dipole attraction between the carbonyl groups

high tensile strength good durability / strong fibre; hydrophobic fibres / water-resistant; chains flexible so fibres easily woven;

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Nylon intermolecular hydrogen bond between –NH...O=C

high specific strength; good durability / strong fibre; resists scratching; resists chemicals/oil; resists UV Stereoisomers(geometric – optical) compounds with the same structural formula but different arrangement of atoms (in space); Geometric isomers Geometrical isomerism in non-cyclic alkenes - C3 and C4 cycloalkanes.

Difference in the physical and chemical properties of geometrical isomers Inter-, intra- molecular H bonds Enantiomers Occur when there are 4 different substituent attached to a single carbon atom (asymmetric carbon atom) Physical properties of enantiomers identical except for rotation of plane polarized light; Chemical properties of enantiomers identical unless they interact with other optically active/chiral compounds/reagents/solvents / identical with achiral compounds/reagents /solvents /

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IB Higher Level Chemical terminology

Term

Definition

Relative atomic mass

The average mass of an atom, taking into account the abundance of all the naturally occurring isotopes, compared to 1/12th of the mass of a carbon-12 atom.

Relative molecular mass

The average mass of a molecule compared to 1/12th of the mass of a carbon-12 atom.

Group

A vertical column of elements on the periodic table with similar chemical properties and the atoms have same valance shell electron configuration.

Period

A horizontal row of elements on the periodic table, across a period from left to right successive atoms have one extra proton and one extra electron in the same outer shell.

Molecular formula

The formula of a molecule that shows the actual number of each type of atom in the molecule.

Empirical formula

The formula of a compound that shows the lowest whole number ratio of each type of atom.

Isotopes

Atoms of the same element with the same number of protons but a different number of neutrons.

Line spectrum

An emission spectrum that has only certain frequencies of visible light that appear as coloured lines.

Ionic bonding

The electrostatic force of attraction between oppositely charged ions arranged in a regular 3D structure called an ionic lattice.

Covalent bond

A shared pair of electrons that holds two atoms together.

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Dative (coordinate) covalent bond

A shared pair of electrons that holds two atoms together and both electrons come from one atom.

Sigma bond

The ‘head on’ over lap of two orbitals along the inter-nuclear axis.

Pi bond

Parallel P-orbitals overlap ‘sideways on’, above and below the inter-nuclear axis.

Hybridization

The mixing of atomic orbitals to form the same number of new orbitals which are equal in energy to one another.

Free radical

A species with one or more unpaired electron.

Resonance

Occurs when one or more valid Lewis structures can be drawn for a molecule or ion and the electrons are delocalized so that they are arranged in an average of the possible Lewis structures.

Delocalisation

A stable bonding arrangement in which electron pairs are not confined to two adjacent bonding atoms but extend over three or more atoms.

Hydrogen bond

A strong intermolecular force of attraction between a delta plus hydrogen and a lone pair of electrons on a very electronegative atom such F, O or N.

Van der Waal’s forces

A weak intermolecular force of attraction that acts between non-polar molecules in the solid and liquid state.

Electronegativity A measure of the ability of an atom to attract a bonding pair of electrons.

Polar bond

A bond in which the electrons are not equally distributed between the two atoms in the bond; the more electronegative atom will have the greater electron density.

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Ligand

A species that uses a lone pair of electrons to form a dative covalent bond with a central metal ion to form a complex.

Periodicity

Repeating pattern of physical and/or chemical properties of elements or their compounds as a function of their arrangement in the periodic table.

Endothermic reaction

A reaction in which heat energy is absorbed from the surroundings causing the temperature to drop.

Exothermic reaction

A reaction in which heat energy is released to the surroundings causing the temperature to increase.

Standard enthalpy of reaction

The enthalpy change of a reaction using the amounts of reactants in a specified stoichiometric equation under standard conditions (298K, 101.3KPa).

Standard state

A reference state for a particular substance under standard thermodynamic conditions (298K, 101.3KPa, 1moldm-3)

Standard Enthalpy of formation, â&#x2C6;&#x2020;HfѲ

The energy change when one mole of compound is formed from its elements in their standard states under standard conditions (298K and 101.3KPa).

Standard enthalpy of combustion

The energy released when one mole of a compound is completely burned in excess oxygen under standard conditions with no change in pressure.

Bond enthalpy

The energy needed to break 1 mol of a specific covalent bond between two atoms in a molecule in the gaseous state averaged over several similar compounds.

First ionization energy

The energy needed to remove one mole of the highest energy electrons from one mole of neutral atoms in the gaseous state under standard conditions (298K, 101.3KPa).

Lattice enthalpy

Energy released when1 mole of a solid ionic compound is formed from its constituent ions in the gaseous state under standard conditions (298K, 101.3KPa).

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Enthalpy of electron affinity

The energy released when one mole of electrons are added to one mole of atoms in the gaseous state under standard conditions (298K, 101.3KPa).

Activation energy

The minimum energy needed for the reactants to react upon collision / The difference in energy between the reactants and transition state.

Catalyst

A substance that increases the rate of reaction by providing an alternative reaction pathway with a lower activation energy and it is not consumed in the overall reaction.

Hetrogenous

Usually applied to a catalyst; in a state or phase different to the reactants.

Overall order of reaction

The sum of the powers of concentration terms in the rate expression.

Rate The slowest step of a reaction. determining step Molecularity

The number of particles in the slowest, rate determining step of the reaction / The number of particles participating in an elementary step of the mechanism.

Half life

The time taken for the concentration of a reagent to decrease to half its original value.

Chemical equilibrium

A chemical reaction is in equilibrium when the rate of the forward reaction is equal to the rate of the reverse reaction and the concentrations of the reactants and products do not change.

Equilibrium constant, kc

The value obtained when the equilibrium concentration of the products are multiplied together and divided by the equilibrium concentration of the reactants multiplied together and all concentrations are raised to the power of their stoichiometric coefficients from the balanced equation.

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Bronsted-Lowry acid Lewis acid Bronsted-Lowry base

A specie that donates a proton.

A specie that accepts an electron pair.

A specie that accepts a proton.

Lewis base

A specie that donates an electron pair.

Weak acid

An acid that partially dissociates in solution.

Strong acid

An acid that fully dissociates in solution.

Conjugate acidbase

A pair of species whose formulae differ by H+.

Buffer solution

A solution that resists change in pH when small amounts of acid or alkali are added, usually made from a weak acid and a salt of its conjugate base or a weak base and a salt of its conjugate acid.

Monoprotic acid

A specie with one replaceable hydrogen atom per molecule.

Diprotic acid

A specie with two replaceable hydrogen atoms per molecule.

Equivalence point

The point in an acid-base titration when the acid and base are present in stoichiometric amounts and all the acid and base has been reacted; it is not necessarily at pH 7.

End point

When an indicator changes colour suddenly during a titration and both the acidic and basic forms of the indicator are present in equal concentration.

Oxidation

When a substance combines with oxygen / When a substance loses hydrogen / When a specie loses electrons / When the oxidation number of an element increases.

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Reducing agent

A species that causes reduction by donating (an) electron(s) to other species and is itself oxidized.

Disproportionation

The simultaneous oxidation and reduction of atoms of a single chemical element in a reaction.

Standard electrode potential Hydrocarbon Saturated

The potential difference measured under standard conditions (298K, 101.3KPa, 1moldm-3 solutions) when a half cell is connected to a standard hydrogen electrode.

A compound containing hydrogen and carbon only.

Molecules that contain only carbon-carbon single bonds.

Unsaturated

Molecules that contain at least one carbon-carbon multiple bond.

Dehydration

The removal of hydrogen and oxygen from a compound in a 2:1 atomic ratio, respectively.

Homologus series

A family of compounds with the same general formula where successive members differ by one CH2 group.

Nucleophilic substitution

The replacement of an atom or group within a molecule, by a species with a lone pair of electrons, which is attracted to an electron deficient part of the molecule.

Isomers

Forms of a compound with the same molecular formula but different structural formulae.

Stereoisomerism Compounds with the same molecular formula and structural formula but different arrangements of atoms in space. There are two types of stereoisomerism; geometrical and optical.

Geometrical isomerism

A type of isomerism that occurs when particular atoms or groups are joined to atoms at each side of a bond that has restricted rotation.

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Optical isomerism

A type of isomerism that occurs when a molecule has no plane of symmetry and can exist in left and right handed forms that are non-superimposable mirror images of each other.

Enantiomer

A molecule that has no plane of symmetry and is nonsuperimposable on its mirror image.

Condensation reaction

When two small molecules combine to form a larger one with the elimination of a smaller molecule (such as water).

Addition polymerisation

Process in which a large number of unsaturated monomers combine to form a polymer without the elimination of any atoms or molecules.

Nucleophile

Negatively charged or neutral species containing an electron pair that is attracted to and can form a covalent bond to electron deficient (delta plus) atom in another molecule.

Functional group An atom or group of atoms responsible for the characteristic reactions of the molecule or homologous series.

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IB HL Chemistry mathematical relationships

Description To calculate the number of moles. Relative molecular mass and relative atomic mass, are determined using the periodic table.

Equation Moles =

To calculate the number of specified particles in a substance. Avogadro’s number NA = 6.02 x 1023 is given in the data booklet.

Number of particles = Moles x Avogadro’s number

To calculate molar concentrations (molarity).

Concentration(moldm-3) = Moles / Volume(dm3)

To convert volumes from cm3 into dm3 for use in calculations.

Volume (dm3) = Volume (cm3) / 1000

Used for dilution calculations. Units can vary but should be used consistently in the equation.

Initial concentration x Initial volume = Final concentration x Final volume

Used to calculate the moles of gas under standard conditions (298K, 101.3KPa)

Moles = Volume(dm3) x 22.4dm3mol-1

The temperature dependent equilibrium constant for the reaction: mA + nB ⇄ xC + yD

Kc = [C]x[D]Y/[A]m[B]n

To calculate the standard entropy change of a reaction from the standard entropies of the reactants and products.

∆SrxnӨ = ΣS(products)Ө - ΣS(reactants)Ө

To calculate the standard enthalpy change of a reaction from the standard enthalpy of formation of the reactants and products.

∆HrxnӨ = Σ∆Hf(products)Ө - Σ∆Hf(reactants)Ө

To calculate the heat energy transferred to change a particular mass of material by a particular temperature.

Heat energy(J) = Mass(g) x Specific heat capacity(JK-1g-1) x Temperature change(K/oC)

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To calculate the charge transferred in a certain amount of time.

Charge(C) = Current(A) x Time(s)

To calculate the moles of electrons transferred in electrolysis. Avogadro’s number and the charge of an electron are given in the data booklet.

Moles of electrons = Charge(C) / Charge of an electron(C) x Avogardo’s number

To work out the Ka (acid dissociation constant) of a weak acid HA. The higher the Ka the stronger the acid.

Ka = [H+][A-] / [HA]

Used to calculate Ka from Kb or vice versa. Kw = 1x10-14 (298K, 101.3KPa)

Ka x Kb = Kw

To calculate pKa from pKb and vice versa

pKa + pKb = 14

Used to calculate pH from pOH and vice versa.

pH + pOH = 14

Used to calculate pH from [H+] and vice versa:

pH = -log10[H+] and [H+] = -10x(pH)

Used to calculate pOH from [OH-] and vice versa

pOH = -log10[OH-] and [OH-] = 10x(pOH)

The Henderson-Hasslebach equation is used to calculate the pH of a buffer solution.

pH = pKa – log10 [Salt]/[weak acid or base]

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Equations in the data booklet

From data booklet

Expanded Pressure(Nm-2) x Volume(m3) = Moles x Molar gas constant(JK-1mol1) x Temperature(K)

Gibbs free energy change (KJmol-1) = Enthalpy change(KJmol-1) – Temperature(K) x Change in entropy(JK-1mol-1) Energy(J) = Plank constant(Js) x Frequency(Hz) Half life(s) = 0.693 / Decay constant(s-1)

Rate constant = Preexponential factor x base of natural logs-(activation energy/Molar

Common use Used to calculate the pressure, volume, temperature or number of moles of an ideal gas.

Used to calculate the Gibbs free energy change. If the value is negative the reaction will be spontaneous under those conditions.

Used to calculate the energy of a single photon from its frequency.

Used to calculate the half life of a reaction or radioactive isotope. First order reactions have a constant half life.

The Arrhenius equation is used to calculate the rate constant k and predict how it will change with activation energy and temperature.

gas constant x Temperature)

Ln(rate constant) = (Activation energy/Molar gas constant x Absolute temperature) + Ln(Preexponetial factor) Log10(Initial intensity / Final intensity) = Molar absorption coefficient (dm3mol-1cm-1) x Path length(cm) x Concentration(moldm-3)

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The logarithmic form of the Arrhenius equation. When LnK is plotted against 1/T a straight line graph is produced with a gradient = -Ea / R Beer-Lambert’s law is used in colorimetry to calculate concentration of a solution when its absorption coefficient is known.

Cells normally have a path length of 1cm and absorbance is measured by the colorimeter. (Log10 Io/I is the absorbance)

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