General Chemical Equations:
Types of oxidation:
1. Metal + oxygen → metal oxide (corrosion)
2. Combustion
I. complete combustion: fuel+ oxygen → carbon dioxide + water (combustion)
II. Incomplete combustion: Fuel + oxygen → Carbon Monoxide (CO) + Carbon +Water
Reactions of acids/metals
3. Metal + acid → a salt + hydrogen
-
2Na + 2HCl → 2NaCl + H₂
4. Alkali metal + water → metal hydroxide + hydrogen
-
2Li + 2H₂O → 2LiOH + H₂
Acid/ Base reactions
5. Metal oxide + acid → a salt + water
-
CuO + 2HCl → CuCl₂ + H₂O
6. Metal hydroxide + acid → a salt + water
-
NaOH + HCl → NaCl + H₂O
7. Metal carbonate + acid → A salt + water + carbon dioxide
-
Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂
Thermal decomposition
8. Metal carbonate → metal oxide + carbon dioxide
-
CaCO₃ → CaO + CO₂.
Basics of Chemistry:
Chemical Reactions: The formation of new substances
Ion: Charged particle
Atom: Smallest possible piece of an element that can exist with the same properties as that
element
Volatile: Substances that can evaporate into gas easily
Mixture: 2 or more substances mingled together but not chemically combined
Molecule: 2 or more atoms/particles covalently bonded together
A: melting
B: Boiling/evaporating
C: Condensation
D: Solidification/Freezing
E: Sublimation
Chapter Checklist:
Ch2: The Atom
✓
Ch14 Oxidation and Reduction:
✓
Ch3: The Arrangement of electrons
✓
Ch15: Volumetric Analysis: Redox
✓
Ch4: The Periodic Table
✓
Ch16: Rates of Reaction
✓
Ch5 Bonding:
✓
Ch17: Chemical Equilibrium
✓
Ch6 Tests for Anions:
✓
Ch18: pH and Indicators
✓
Ch7: Trends in Periodic Table
Ch19: Environment Chemistry - Water
Ch8: Radioactivity
✓
Ch20: Electrochemistry
✓
Ch9: The Mole Concept
✓
Ch21: Fuels and Heats of Reaction
✓
Ch10: Properties of Gases
✓
Ch22: Some families of Organic Compounds
Ch11:Stoiciometry I
✓
Ch23: Types of Reactions in Organic Chemistry
Ch12 Acids and Bases:
✓
Ch24: Stoichiometry II
Ch13 Volumetric Analysis:
✓
Experiments:
1.
Ch3: Flame tests
✓
15. Ch21: b) Preparation and properties of ethyne
✓
2.
Ch6: Anion Test (a-e)
✓
16. Ch21: c) Determination of the heat of reaction
of hydrochloric acid and Sodium hydroxide
✓
3.
Ch9: To measure the relative molecular mass
of a volatile liquid
✓
17. Ch14: Oxidation-Reduction reactions:
A. Halogens as oxidising agents
B. Displacement reactions of metals
✓
4.
Ch13: a) To prepare a standard solution of
Anhydrous Sodium Carbonate
✓
18. Ch17: Equilibrium experiments to illustrate Le
Chatlier’s principle
5.
Ch13: b) To use a standard solution of sodium
carbonate to standardise a given hydrochloric
acid solution
✓
19. Ch19: a) Determination of the concentration of
free chlorine in swimming pool water or bleach
6.
Ch13: c) To determine the concentration of
ethanoic acid in vinegar
✓
20. Ch19: b) Determination in water of total i)
suspended solids, ii) total dissolved solids and
(iii) pH
7.
Ch13: d) To determine the amount of water
crystallisation in hydrated sodium carbonate
✓
21. Ch19: c) Estimation of total hardness of a
water sample
8.
Ch15: a) To prepare a standard solution of
Ammonium Iron (||) Sulfate to standardise a
solution of potassium permanganate
✓
22. Ch19: d) Estimation of dissolved oxygen by
redox titration
9.
Ch15: b)To determine the amount of iron in an
iron tablet
✓
23. Ch?:
A. To extract clove oil from cloves by steam
distillation
B. To isolate eugenol from an emulsion of
clove oil and water by solvent extraction
10. Ch15: c)To prepare a solution of Sodium
Thiosulfate and to standardise it by titration
against a solution of iodine
✓
24. Ch18: To prepare a sample of soap
11. Ch15: d)To determine the percentage (w/v) of
Sodium Hypochlorite in bleach
✓
25. Ch22: a)
A. Properties of ethanol
B. Properties of ethanoic acid
12. Ch16: a) Measuring the rate of the production
of oxygen from hydrogen Peroxide
✓
-
26. Ch22: b) Preparation of Benzoic acid from
phenyl methanol
13. Ch16: b) To study the effect on reaction rate
of concentration and temperature
-Disapperaing cross
✓
-
27. Ch18: To separate a mixture of indicators or
coloured substances Using paper/ thin layer/
column chromatography
14. Ch21: a) Preparation and properties of ethane
✓
28. Ch?: Recrystallise Benzoic Acid and determine
its melting point
Ch2: The Atom
Particulate nature of Matter:The idea all matter is made up of very small particles
-
Demonstration: Concentrated Ammonia 𝑁𝐻3 and HCl soaked in cotton in a tube →cloudy
ring of ammonium chloride
Diffusion:The spreading of particles (in a gas or liquid) from a high area of concentration to an
area of lower conc
Dalton’s Theory: all matter is made up of tiny particles called atoms . Atoms are indivisible
Cathode ray: Stream of electrons / Travel Cathode to Anode- Is able to fluoresce glass- Is strong
enough to move a paddle-can be deflected/attracted by mag. fields
Cathode: Electrode connected to the neg. end of a battery
Alpha particle: a helium nucleus emitted by some radioactive substances,
Proton: a subatomic particle found in the nucleus with a positive charge
Neutron: a subatomic particle of about the same mass as a proton but without a charge
Electron: a subatomic particle with a negative charge
Nucleus: Dense core within an atom containing the protons and neutrons
Subatomic particle: a small portion of matter within the Atom- Protons/Neutrons/Electrons
Subatomic: Any particles of matter that are smaller than a hydrogen atom
Particle: A particle is a small portion of matter
Thomson’s model of the Atom: Plum pudding model/ A sphere of positive charge, electrons
randomly embedded throughout.
Discovery Timeline:
Greek Philosophers/Democritus: Came up with the concepts of small indivisible particles called
atoms
John Dalton: Wrote books proposing the idea of the atom - started looking for experimental
evidence
William Crookes: Investigated cathode rays in vacuum tubes
George Stoney: Proposed the name Electron
−
𝑒
J.J. Thompson: Showed 𝑒 are negative and found the 𝑚 of the electron (ratio bet. size of
charge/mass) His research on cathode rays led him to discover the electron.
Robert Millikan: Oil drop experiment (air was ionised/ lost electrons-electrons picked up by oil
droplets) found the magnitude of the charge on electron
-Thompson’s plum pudding model
Ernest Rutherford: discovered the nucleus with Gold foil exp/and the proton by bombarding
lighter elements with alpha particles.
James Chadwick: Discovered the neutron Radioactive source-> released Alpha particles-> Beryllium->released Neutrons->Paraffin wax
block->released protons - allowed neutrons to be measured
Uses of Neutrons: splitting of uranium atoms to produce nuclear energy in nuclear rectors and
atomic bombs
Rutherford's prediction: small deflections or go straight through since he believed that + charge
was spread evenly and not strong
Observation
Conclusion
Most alpha particles pass straight through the
foil
Most of the atom is empty space
Some alpha particles are deflected at large
angles
Alpha particles are repelled because they pass
a small positive nucleus
A small number of alpha particles are
deflected along their own path
The alpha particles had collided head on with
the nucleus - the nucleus is very small
Ch3: The Arrangement of Electrons
Spectrum: A spectrum is a set of wavelengths of light/spread of colours
Prism: a transparent solid object often having triangular bases, used for dispersing light into a
spectrum or for reflecting rays of light
Continuous Spectrum: a continuous spread of colours ROYGBIV- formed when white light is
passed through a prism
Line spectrum: coloured lines against a dark background formed when light emitted from a
discharge tube filled with gas passes through a prism
Atomic absorption spectrum: consists of dark lines along a coloured background formed when
white light passing through a gaseous sample strikes a prism
Uses- To identify and find the concentrations of elements in gases
Spectroscope: Instrument used to observe Spectra
Spectrometer: Instrument used to measure and observe Spectra
Bohr’s theory:
●
●
●
●
●
Electrons have a set amount of energy
They travel in fixed orbits a set distance away from nucleuse- of diff energy travel along diff paths.
E2-E1= h(Plack’s constant) x f ( light emitted)
-showed that n levels become closer the higher the energy they are
-Quantisation: the idea that electrons in an atom may have specific energy values- each energy
lvl has diff E.V. / travels in a fixed paths
-Limitations: Failed to take into account sublevels/wave particle-duality
Quantum: A fixed amount of energy an atom may have
Louis De Broglie (wave-particle duality)-electrons move in a wave motion - disproved Bohr’s idea
that electrons are at fixed distances from the nucleus
Hesienberg's Uncertainty Principle:it is impossible to measure at the same time both the
velocity and position of an electron
Schrodinger: worked out the probability of finding a particular electron in an atom.
Energy level: A fixed energy value that an electron in an atom may have in an atom
Ground state: is when the electrons in an atom occupy the lowest available energy level.
Excited state: energy level of an atom, ion, or molecule in which an electron is at a higher energy
level than its ground state. Is unstable so doesn't last
Balmer series: Lines of the visible line spectrum of Hydrogen (drops down from higher energy
levels to n=2)
N3→n2 red
Paschen series (IR): invisible lines of the line spectrum of Hydrogen (seen with InfraRed)
Lyman(UV): invisible lines of the line spectrum of Hydrogen (seen with UltraViolet)
Sublevel: A subdivision of an energy level consisting of one or more orbitals of the same energy
Orbital: A region in space around the nucleus of an atom where there is a high probability of
finding an electron
Orbit: is the definite path of an electron that moves around the nucleus in an atom
S-orbital: spherical shape
P-orbital: Dumbbell-shaped, found along 1 of 3 axes, which are at right angles to each other
Less attraction to outer shell → more reactive:
Salts and their Coloured flames:
Metal Present
Colour
Lithium
Crimson
Potassium
Lilac
Sodium
yellow
Strontium
Red
Copper
Blue-Green
Barium
Green
●
●
●
●
●
Light a bunsen burner and adjust it to obtain a blue flame.
Dip the wooden splint in water and add one of the salt to it so it sticks on.
Place the wooden splint over the flame to observe the colour of the flame.
After the salt has finished burning, put the flame out and take note of the colour.
Repeat this for all the different salts.
Ch4: The Periodic Table
Element: Substance that cannot be divided into simpler substances by chemical means/ Pure
substance made up of one type of atom
Robert Boyle: defined the element as a substance that cannot be divided into simpler
substances by chemical means
Humphrey davy: Discovered many elements using electricity
Johann Dobereiner: Classified Elements into triads
Triads: Group of elements with similar chemical properties in which the Atomics weight of the
middle element is the approx. average of the other two elements
John Newlands/Newland’s octaves: Arrangement of elements in which the first and eight
elements, counting from any particular element, have similar properties.
*Flaws
➢ No gaps
➢ Some properties in the same group were completely different
Mendeleev’s periodic Law: when elements are arranged in order of increasing atomic weight,
the properties of the elements recur periodically
*Difference between Octaves and Mendeleev’s laws:
❖ Gaps in Mend.PT - none in octaves.
❖ Elements in the same group always have similar properties.
*Difference between Modern Periodic Table and Mendeleev’s table:
❖ Gaps in Mend.PT - No gaps in Mod P.T.
❖ Ordered in increasing atomic weight(Mend P.T.)/ increasing atomic number (Mod P.T.)
❖ No noble gases (Mend PT)/ Noble gases present (Mod PT)
The Periodic Law: when elements are arranged in order of increasing atomic number, the
properties of the elements recur periodically
Henry Moseley: Used X-rays to determine Atomic no. in elements/corrected Mend’s table
Processes in the Mass Spectrometer:
➔ Vaporisation:Liquids are vaporised into gas vacuum
Ionisation: Electrons are knocked out of the gas particles by an electron gun making them
+ly charged
Acceleration: The particles are accelerated to high velocities by neg. charged plates
Separation in Mag. field: mag. Fields deflect the beam of particles in accordance to their
mass
Detection: is able to record the beam- m.f. Can be changed to measure diff. beams
Principle of mass spectrometry: Charged particles are deflected at different extents according
to their mass
-Uses: Detects the presence/relative abundance of isotopes/ Identify compounds e.g. Drug
testing
Mass Spectrum: Graph showing rel. atomic mass and rel. Abundance of isotopes
Electron configuration: A description of how electrons are laid out in an atom’s energy levels
Aufbau Principle: when building up the electronic configuration of an atom in its ground state,
the electrons occupy the lowest available energy level
Pauli’s Exclusion Principle: no more than two electrons can occupy an orbital and they must
have opposite spin
Hund’s Rule: when two or more orbitals of equal energy are available, electrons occupy them
singly before filling them in pairs.
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, compared with 1/12th of the mass of
the carbon-12 isotope
Mass number: The sum of the protons and neutrons in the nucleus of an atom
Atomic number:The number of protons (or electrons) in a neutral atom
Isotope: are atoms of the same element (i.e. same atomic number) that have different mass
numbers due to the different number of neutrons in the nucleus
Ch10: Properties of Gases
Solid: A substance which has a fixed volume, fixed shape and whose particles can only vibrate in
a fixed position
Liquid: A substance that has a fixed volume and fills the base of the container it is in. Its particles
can vibrate, translate and rotate
Gas: A substance that has no well-defined boundaries but diffuses rapidly to fill any container in
which it is placed- particles rotate, vibrate, translate and diffuse
Temperature: a measure of the hotness of an objects
Pressure: The amount of pressure a gas exerts on the container it's in
Volume: The amount of space a gas fill up
Boyle's law: at constant temperature the volume of a fixed mass of gas is inversely proportional
to its pressure V ∝ 1/P → VxP=K
Charles' law: at a constant pressure the volume of a fixed mass of gas is directly proportional to
its temperature measured on the kelvin scale V ∝ T → V/T=K
Combined gas law: it states that the ratio of the product of pressure and volume over
temperature of a gas is equal to a constant.
Gay-Lussac's law: in a reaction between gases the volume of the reacting gases and the volume
of any gaseous products are in the ratio of small whole number while at the same temperature
and pressure
Avogadro's law: equal volumes of gases contain equal number of particles under the same
conditions of temperature and pressure- Volume is proportional to number of Particles
Brownian motion: random movement of particles suspended in a liquid or gas
Kinetic theory:
●
●
●
●
Gases have rapid, continuous, random movement of particles which collide with each
other
They have no attractive/repulsive forces
Volume and size of each particle is negligible as they are small
Collision are elastic -energy is transferred not lost between the particles
Ideal gas: is one that obeys all the assumptions of the kinetic theory of gases under all
conditions of temperature and pressure
Real gas: Gases are not Ideal gases since a) repulsion/attraction bet. molecules exist at low
temp/high pressure b) the volume of a molecule is not negligible at high pressure
Ch11:Stoiciometry
Molecular formula: formula that gives the actual number of each atom in the molecule
Empirical formula: formula which gives the simplest whole-number ratio of the elements in the
compound
Gravimetric analysis : Calculations of chemical reactions based on their masses
Mass: a measure of the amount of matter that an object contains
Law of Conservation of Mass: total mass of the reactants are equal to the total mass of the
products of a chemical reaction
Law of Conservation of Matter: in any chemical reaction matter is neither created nor
destroyed but merely changed from one form to another
Ch5: Bonding
Compound: is a substance that is made up of two or more different elements chemically
combined
Octet rule: states when bonding occurs, atoms tend to reach an electron arrangement with eight
electrons in the outermost shell
Ionic bonding: The electrostatic attraction between positive and negative ions formed by the
transfer of electrons.. This forms a three-dimensional lattice.
Covalent Bonding: The sharing of a pair of electrons.
-often non-metals/generally small/gases at room temperature.
-Giant covalent molecules e.g diamond and graphite (allotropes of carbon).
Lattice: repeated units/patterns
Crystal Lattice: 3d arrangement of ions
Transition metal: is one that forms at least one ion with a partially filled d sublevel
Valency: the number of atoms of hydrogen or any other monovalent element with which each
atom of that element combines
VSEPR Theory: Valencey, Shell, Electron, Pair, Repulsion Theory.
Electronegativity: the relative attraction that an atom in a molecule has for the shared pair of
electrons in a covalent bond
-Elements with low electronegative are referred to as electropositive +will carry a slight positive
charge in a bond.
Lone Pairs: not involved in bonding (opposite bond pairs)
Atomic radius: :half the distance between the nuclei of two atoms of the same element that are
joined together by a single covalent bond
First Ionisation energy: (of an element) is the energy required to remove the most loosely
bound electron from a neutral gaseous atom in its ground state
Sigma bond (σ): overlap of orbitals - Pi bonds(π)- overlap of p orbitals
Single bond: a shared pair of electrons between two atoms (two overlapping s orbitals or two
head-to-head p-orbitals)
Dative covalent bond: Both electrons forming the covalent bond, have come from the same
atom.
Double bonds- one sigma one pi
triple bonds- one sigma 2 pi
Diff between Ionic/Covalent Compounds:
●
●
●
●
Hardness- Ionic Hard - Covalent - molecular crystals soft
Melting/boiling points: Ionic high melting point/ Covalent low melting point
Conduction of electricity: ionic yes if melted dissolved in water / covalent no
Ionic contains a network of ions in a crystal / covalent= individual molecules
Lone pairs affect this- repels elements
Electronegativity difference more than 1.7
E.Neg diff equal/less than 1.7
E. Neg diff less than/equal to 0.4
E.Neg. diff between 0.4 or 1.7
- ionic
- covalent
- nonpolar covalent
-polar covalent
Polar covalent bond: an unequal sharing of the electron pair. This causes the molecule to be
delta positive δ + and delta negative δ - .
*Polar bonds do not always mean the molecule is polar. If the molecule is symmetrical, bond
polarity can be cancelled out.
Pure covalent: equal sharing of a pair of electrons. Only found in molecules of the same
element.
*Mabelle-> malleable
Intramolecular bonding(bonds between atoms in a molecule)
i. Pure covalent
ii. Polar covalent
Intermolecular bonding(bond between molecules
i. Van der Waals
ii. Dipole-dipole
iii. Hydrogen bonding
Van der Waals forces: The weakest intermolecular force which occurs in pure covalent and
nonpolar covalents . It is caused by electrons in a molecule spontaneously being unevenly
distributed within the molecule, causing a slight positive and negative pole to be produced
momentarily. This then induces opposite poles in neighbouring molecules. Oppositely charged by
poles attract.
Dipole-dipole forces: This is the slight attraction between oppositely charged ends of a polar
molecule.
Hydrogen bonding: Weak intermolecular force between molecules caused byalternating
temporary bonds between the electropositive hydrogen and a lone pair of electrons of other
molecules This occurs in molecules containing a Hydrogen atom and either Oxygen, Nitrogen or
Fluorine atoms.
Ch8: Radioactivity
Henri Bacquerel: Discovered Radioactivity
-When some Uranium salt was wrapped in blackpaper, it fogged the photographic plate it was
lying on
-found out that radiation was occurring and it was spontaneous- occurring naturally
The Curies: Pierre and Marie Curies investigated radiation
-Discovered: No.84 Polonium
No.88 Radium
-1903 Becquerel and the Curies won the Nobel Prize in Physics
-1911 Marie Curie won a Nobel Prize in Chemistry
Radioactivity: The spontaneous breaking up of unstable nuclei with the emission of one or more
types of radiation.
-
Types:
-Alpha particles
-Beta particles
-Gamma particles
Alpha Particles α: A group of 2 protons and 2 neutrons stuck together (a helium nucleus)
emitted by certain radioactive substances
●
●
●
●
double positive charge
cannot penetrate a few cm of air/a sheet of paper
is slower than other types of radiation due to relatively large mass
isn’t hazardous- Americium-241 is used in smoke detectors
Beta Particles β : Fast moving electron emitted when a neutron in an unstable nucleus breaks
down into a proton
●
●
●
is faster due to smaller size
can penetrate up to 5mm of Aluminium
Carbon-14 is an e.g. - used to find age of objects
●
●
●
●
Gamma Radiation γ: High energy Electromagnetic waves emitted by an unstable nucleus
to remove surplus energy
moves at the speed of light
is not a charged particle- isn’t deflected by electric or magnetic fields
High penetration- can only be stopped by thick slab of lead
Dangerous- can alter structure of chemicals within the body
-> Higher risk of Cancer/Can kill Cancer cells
e.g. Cobalt-60 - food irradiation
●
Geiger-Muller tube: used to detect radiation
-ratemeter - measures no. of radioactive particles entering per sec
-loudspeaker- indicates presence of activity
Background count/background radiation : a measure of natural radioactivity
E.g. cosmic rays from the sun, from rocks
Nuclear reactions: The altering of the composition, structure and energy of an atomic
nucleus
Chemical Reactions
Nuclear reactions
Involves changes in electron shell
Involves changes in the nucleus
No new element formed
New element formed
No nuclear radiation
Nuclear radiation
Chemical bonds broken/formed
No formation/breaking of chem. bonds
Half Life: Time taken for half of a samples' unstable nucleus to decay
Radioisotopes: Isotopes with an usable nucleus which emits radiation
Uses:
Medicine:
-Radiotherapy- using Gamma rays to kill cancer cells and germs
-Iodine-131- beta measures uptake of iodine by the thyroid gland
Archaeology:
-Carbon Dating: finding the age of objects by finding ratio of Carbon-12 Carbon-14
-Once an organism dies- carbon-14 stops being replaced and decays (half life = 5730)
Agricultural:
-Phosphorus-32 used in fertilisers Carbon-14 helps with fertilisers
Irradiation:
-Co-60-using gamma rays to kill bacteria on food
Smoke Alarms:
Industry: Locate leaks
Alpha Particle decay:
Beta Particle decay:
Gamma Radiation:
X → X + gamma radiation
Important Ions:
Name
Formula
Hydroxide ion
O𝐻
Nitrate ion
𝑁𝑂3
Charge of the ion
−
−
One negative charge
Hydrogencarbonate ion
−
𝐻𝐶𝑂3
−
Permanganate ion
𝑀𝑛𝑂4
Carbonate ion
2−
𝐶𝑂3
2−
Chromate ion
𝐶𝑟𝑂4
Dichromate ion
𝐶𝑟2𝑂
2−
7
Two negative charges
2−
Sulfate ion
𝑆𝑂4
Sulfite ion
𝑆𝑂3
Thiosulfate ion
𝑆2𝑂
2−
2−
3
Phosphate ion
3−
Three negative charges
+
One positive charge
𝑃𝑂4
Ammonium ion
𝑁𝐻4
Ch6: Tests for Anions
a) Test for Chlorine (aq):
●
Test: Add drops of Ag𝑁𝑂3 to an aqueous solution.
●
Observation:
Clear to Cloudy + White precipitate is formed
Is soluble in dilute ammonia - cloudiness disappears when added
+
−
𝐴𝑔 + 𝐶𝑙 → AgCl↓
b) Test for Sulfate/Sulfite ions(aq):
●
Test: Add 𝐵𝑎𝐶𝑙2 to a an aqueous solution
Distinguishing test: Add dilute HCl
●
Observation:
Clear to Cloudy + White precipitate formed
Sulfate: precipitate remains
Sulfite: precipitate dissolves
2+
+ 𝑆𝑂4
2+
+ 𝑆𝑂3 →𝐵𝑎𝑆𝑂3 ↓
𝐵𝑎
𝐵𝑎
2−
→𝐵𝑎𝑆𝑂4 ↓
2−
𝐵𝑎𝑆𝑂4 +HCl→ no change
2−
𝑆𝑂3
+ 2H →𝑆𝑂4 + 𝐻2O
c) Test for Carbonate/Hydrogen-carbonate ions(aq):
●
Test: Salt placed in tilted boiling tube + dilute HCl is added
Stopper is placed and tubing is placed into a test tube of limewater
Distinguishing test: Add 𝑀𝑔𝑆𝑂4 to fresh aqueous sol of the salts
●
Observation:
Fizzing + 𝐶𝑂2 produced + limewater becomes cloudy in appearance
Carbonate : white precipitate formed
Hydrogen Carbonate: no precipitate formed unless boiled
2−
+
𝐶𝑂3 +2𝐻 →𝐶𝑂2 + 𝐻2O
−
+
𝐻𝐶𝑂3 +𝐻 →𝐶𝑂2 + 𝐻2O
𝐶𝑎(𝑂𝐻)2 +𝐶𝑂2→𝐶𝑎𝐶𝑂3 ↓ + 𝐻2O
2+
2−
𝑀𝑔 +𝐶𝑂3
2+
→𝑀𝑔𝐶𝑂3 ↓ (insoluble)
−
𝑀𝑔 +2𝐻𝐶𝑂3 →𝑀𝑔(𝐻𝐶𝑂3 )2 (soluble)
𝑀𝑔(𝐻𝐶𝑂3 )2 → 𝑀𝑔𝐶𝑂3 ↓ +𝐶𝑂2 + 𝐻2O
d) Test for Nitrate ions(aq):
●
Test: Add 𝐹𝑒𝑆𝑂4 sol to an aqueous sol of the salt
Slant test tube + add concentrated 𝐻2 𝑆𝑂4 down the inside
●
Observation:
A brown ring is formed at the junction of the two liquids
e) Test for Phosphate (aq):
●
Test: Ammonium molybdate is added to a solution of the salt
Concentrated nitric acid is + solution warmed
●
Observation:
Yellow precipitate is formed
Ch12: Acids and Bases
Arhennius’s definition of
-Acid : A substance which dissociates in water to form hydrogen ions(Protons)
-Base: A substance which dissociates in water to form hydroxide ions(anions)
Limitations:
Limited to acids and bases which dissolve in water
Doesn’t take into account amphiprotic nature of substances
Bronsted-Lowry Theory:
Acid: Proton donor
Base: Proton acceptor
Strong/Weak Acid/Base : Good/bad proton donor/acceptor
Amphiprotic/Amphoteric: Substance which can act as an acid and a base
Monobasic/Monoprotic: Acids with one H+
Conjugate Acid-base pair: A pair consisting of an acid and a base which differ by one proton
Examples of Acids and bases :
Acids:
Ethanoic/Vinegar
Sulfuric Acid/battery acid
Bases:
NaOH (caustic soda) used in oven cleaners
Bleach (Sodium hypochlorite)-oxidising agent
Salt: substance formed when the hydrogen ion from an acid is replaced by a metal or an
ammonium ion
Neutralisation: the reaction between an acid and a base to form a salt and water
1. Medicine:
Excess of HCl (in the stomach) caused by overeating can be harmful
Antacids can be taken to neutralise the acid
E.g Alka-Seltzer and Bisodol contain sodium hydrogen-carbonate
Milk of Magnesia and Maalox contain 𝑀𝑔(𝑂𝐻)2
2. Agriculture:
If soil is too acidic, the yield of crops tends to be low
Lime (CaO) is often added to water to form 𝐶𝑎(𝑂𝐻)2 (slaked lime) which neutralises the
acid + 𝐶𝑎𝐶𝑂3 can also be used (Ground Limestone)
3. Environment protection:
Areas suffering from acid rain have Limestone added - it is also added to acidic coal
chimneys
4. Miscellaneous:
Toothpaste is slightly basic to neutralise acids causing tooth decay
Baking soda can neutralise acidic sting of bees
Vinegar can neutralise the alkaline sting of wasps
Shampoo is slightly basic to open the scales coating hair
Hair conditioner neutralise the shampoo to seal up the scales and leave the hair more
manageable and shiny
Ch14: Oxidation and Reduction:
O.I.L. R.I.G.
Oxidation is loss (of electrons) + an increase in oxidation no.
Reduction is gain (of electrons) + a decrease in oxidation no.
Oxidising agent: Substance which brings about oxidation in another substance i.e. gains
electrons from the substance which has oxidised.
Reducing agent: Substance which brings about reduction in another substance i.e. loses the
electrons which is gained by the other substance
Chlorine - oxidising agent
Hydrogen Peroxide - oxidising agent (used in bleaching hair)
Rules for assigning Oxidation no.:
●
●
●
The oxidation for any uncombined element is zero
The oxidation no of an ion is the same as its charge
The sum of the oxidation numbers of all elements in a compound must add up to
zero
●
●
●
●
Oxygen always has an oxidation number of -2 except in peroxides e.g. 𝐻2 𝑂2 (-1) and
𝑂𝐹2 (+2)
Hydrogen is always assigned +1 in its compounds except in metal hydrides where it
is -1 (NaH)
Halogens always have an oxidation number of -1 in their compounds unless bonded
to a more electronegative atom.
The sum of the oxidation numbers of all elements in a complex ion (compounds
must equal the charge on the ion.
Halogens as oxidising agents :
●
A). Bromine + Chlorine:
● 1 fifth of test tube is filled with Potassium Bromide
● Same amount of chloride water is added
Observation:
●
Potassium Bromide: colourless → yellow/orange
−
−
Oxidation: 2𝐵𝑟 → 𝐵𝑟2 + 2𝑒
−
−
Reduction: 𝐶𝑙2 + 2𝑒 → 2𝐶𝑙
●
B). Iodine + Chlorine:
● 1 fifth of test tube is filled with Potassium iodide
● Same amount of chloride water is added
Observation:
●
Potassium iodide: colourless → reddish-brown
−
−
Oxidation: 2𝐼 → 𝐼2 + 2𝑒
−
−
Reduction: 𝐶𝑙2 + 2𝑒 → 2𝐶𝑙
●
C). Iodine + Bromine:
● 1 fifth of test tube is filled with Potassium iodide
● Same amount of bromine water is added
Observation:
●
Potassium iodide: colourless → reddish-brown
−
−
Oxidation: 2𝐼 → 𝐼2 + 2𝑒
−
−
Reduction: 𝐵𝑟2 + 2𝑒 → 2𝐵𝑟
●
D). Sodium Sulfate :
● 1 fifth of test tube is filled with Sodium Sulfate
● Same amount of chlorine (or Br/I) water is added
● Barium Chloride added + shaken
●
Dilute HCl added
Observation:
●
●
●
+ Chlorine water : no change
+ Barium Chloride: white precipitate formed
+ HCl: Precipitate does not dissolve
Conclusion:
●
The not dissolving of precipitate proves presence of sulfate
2−
2−
+
−
Oxidation: 𝑆𝑂3 +𝐻2O→𝑆𝑂4 + 2𝐻 + 2𝑒
−
−
Reduction: 𝐶𝑙2 + 2𝑒 → 2𝐶𝑙
●
D). Iron Sulfate :
● 1 fifth of test tube is filled with Iron Sulfate
● Same amount of chlorine (or Br/I) water is added
● Barium Chloride added + shaken
● Dilute NaOH added
Observation:
●
●
+ Chlorine water : no change
+ Sodium hydroxide: greenish/brown precipitate formed (iron oxide)
2+
Oxidation: 2𝐹𝑒
3+
−
→ 2𝐹𝑒 + 2𝑒
−
−
Reduction: 𝐶𝑙2 + 2𝑒 → 2𝐶𝑙
Displacement reaction of metal:
●
E.) Zinc:
● Add Powdered/granulated zinc to copper sulfate solution
Observation:
●
●
Zinc: grey → dark brown
Copper sulfate: blue →colourless
2+
−
Oxidation: Zn → 𝑍𝑛 + 2𝑒
2+
−
Reduction: 𝐶𝑢 + 2𝑒 → Cu
●
E.) Magnesium:
○ Sand magnesium ribbon with sandpaper
○ Dip magnesium into copper sulfate solution using tongs
Observation:
●
●
Magnesium: → dark brown coat
Copper sulfate: blue →colourless
2+
−
Oxidation: Mg→ 𝑀𝑔 + 2𝑒
2+
−
Reduction: 𝐶𝑢 + 2𝑒 → Cu
Ch15: Volumetric Analysis: Oxidation-Reduction
Potassium permanganate(potassium manganate(VII)): 𝐾𝑀𝑛𝑂4
-Purple solid /not a primary standard
-cannot be obtained in a state of high purity / decomposes in the presence of
sunlight+heat
- has to be standardised by titration against a primary standard solution.
Reaction in an acidic solution:
(+7)
(+2)
−
+
−
2+
𝑀𝑛𝑂4 +8𝐻 + 5𝑒 →𝑀𝑛
+ 4𝐻2O
Purple —-------------> colourless
+
-Dilute acid is needed supply 𝐻 ions - sulfuric acid is used HCl can’t - oxidises the ions Nitric acid - strong oxidising agent + interferes with the reaction
Reaction in an basic solution:
(+7)
(+4)
−
−
−
𝑀𝑛𝑂4 +2𝐻2𝑂 +3𝑒 →𝑀𝑛𝑂2 +4𝑂𝐻
Purple —-------------> brown
●
●
Potassium permanganate is self-indicating
Reading is taken from top of the meniscus
−
2+
Reaction of 𝑀𝑛𝑂4 and 𝐹𝑒 :
−
2+
𝑀𝑛𝑂4 is often titrated against 𝐹𝑒
−
+
2+
𝑀𝑛𝑂4 +8𝐻 + 5F𝑒
●
●
●
●
●
●
●
2+
→𝑀𝑛
3+
+5𝐹𝑒
4𝐻2O
Mn:Fe 1:5
2+
𝐹𝑒 is not obtained by dissolving crystal iron (II) sulfate 𝐹𝑒𝑆𝑂4 .7𝐻2O
- crystals are oxidised by air
-Lose water crystallisation when exposed to air(efflorescence)
Anhydrous 𝐹𝑒𝑆𝑂4 can’t be used - takes in moisture from the air (deliquescence)
Ammonium iron (II) sulfate/ ferrous Ammonium sulfate used
- unaffected by air / is obtained in a high degree of purity
(𝑁𝐻4 )2𝐹𝑒(𝑆𝑂4 )2.6𝐻2O - high Mr of 392 reduces margin of error
Is used as a primary standard even though it is hydrated as it is stable and doesn’t lose
water of crystallisation to air
Ammonium iron (II) sulfate is dissolved in water + dilute Sulphuric acid- prevents
oxidation by oxyge air
Procedure is in the experiment section:
End point: Purple → Permanent faint pink
Autocatalysis: When the rate of reaction increases as the ions are formed
Reaction between Iodine and Sodium Thiosulfate(𝑁𝑎2𝑆2𝑂3):
●
𝑁𝑎2𝑆2𝑂3 a.k.a. Sodium hyposulfite - reducing agent used in analysing the amount of certain
●
●
substances dissolved in water e.g Oxygen and Chlorine, for extracting gold and as an
antidote to cyanide poisoning
Also used for developing photos before digital cameras were invented
Most commonly found as a colourless crystalline solid 𝑁𝑎2𝑆2𝑂3.5𝐻2O
(+2)
(+2.5)
2−
2−
−
2−
𝐼2 + 2𝑆2𝑂3 → 𝑆4𝑂6 +2𝐼
𝑆4𝑂6 = sodium tetrathionate
Not a primary standard- can’t be sufficiently found in a pure state/ crystals are efflorescent
-Therefore a known solution cannot be made
It is standardised by titrating it against Iodine
Iodine is also not a primary standard - sublimes even at RT and it is almost fully insoluble
in water
A standard solution can be obtained by titrating against Potassium Permanganate
●
●
●
●
−
−
+
2+
2𝑀𝑛𝑂4 +10𝐼 +16𝐻 →2𝑀𝑛
Ratio: Mn:I is 2:5
●
+5𝐼2 +8𝐻2O
Excess potassium iodide is added to form the triiodide ion in the solution to add solubility
and polarity to the non-polar sparingly soluble Iodine and so all the KMnO4 reacts
complete to give a precise measurement of the Iodine
End point: Reddish brown Iodine →becomes paler -> yellow→ colourless
Starch used to help pinpoint the end point.
Colour Change: Blue/black → colourless
Starch is added when the solution becomes pale yellow. Since if it is added earlier it would be
absorbed by the high quantity of Iodine and would make the reaction slower
Determining w/v of sodium hypochlorite (NaClO) in bleach:
To find out conc of NaClO you titrate against an excess potassium iodide and then titrate the
liberated Iodine against Sodium Thiosulfate.
Why is bleach diluted?:
Bleach has a high concentration so would require a large amount of iodine to react with it
Ch16: Rates of Reaction
Rate of reaction: Change in concentration per unit time of any one reactant or product
Ways to measure rate of reaction:
●
●
●
Inverted graduated cylinder (change in vol of products)
Gas syringe (change in vol of products)
Balance (change in mass of reactants)
Instantaneous Rate of reaction: rate of reaction at any one particular time during the reaction
Factors affecting rate of reaction:
●
Nature of reactants:
Ionic compounds react quicker than covalent bonds since covalent bonds take longer to
be broken down and remade
E.g Acidified Sodium dichromate 𝑁𝑎2𝐶𝑟2𝑂7 + (𝑁𝐻4 )2𝐹𝑒(𝑆𝑂4 )2.6𝐻2O (ionic)
Is faster than
𝑁𝑎2𝐶𝑟2𝑂7 + ethanal (covalent)
●
Particle size:
Increases number of collisions / increases number of correct orientation
E.g. Calcium Carbonate (marble chips + HCl)
●
Concentration:
Increases number of collisions / increases number of correct orientation
●
Temperature:
Increased energy / increases number of collisions
●
Catalysts:
Catalyst: a substance that alters the rate of reactions of a chemical reaction but is not
consumed in the reaction
Inhibitors/ negative catalyst: Catalysts which slow down reactions
a.) General properties:
■ Are recovered chemically unchanged at the end of a reaction
■ They tend to be specific (each type reacts with different substances)
■ Enzymes: substances that is produced by a living organism to be used as a
biological catalyst
E.g. Catalase + 𝐻202
■ Need to be present in only small quantities
■ Helps equilibrium to be achieved
■ Catalyst poisons prevent catalysts from working e.g Lead vs catalytic
converters
b.) Types of catalysts:
■ Homogeneous catalysis: Reactants and catalyst are in the same phase (no
boundary between them)
■ Heterogeneous catalysis: Reactants and catalysts are in a different phase
■ Autocatalysis: Catalysis in which one of the products acts as a catalyst
c.) Mechanism of catalysts:
■ 1. The Intermediate Formation Theory:
the catalyst forms an unstable intermediate compound with the reactants
and this unstable intermediate is then decomposed to form the desired
products.
■ 2. The Surface Adsorption Theory (Heterogeneous Catalyst):
Adsorption - accumulation of substances on the surface of a substance
● Step1 Adsorption occurs using temporary bonds to bring reactants
closer together and weaken covalent bonds
● Step2 Reaction
● Step3 Desorption stage: reacted reactants leave and get replaced
d.) Catalytic converters:
○ Catalytic converter: a device in the exhaust system
containing catalysts which convert pollutants in exhaust gases
to less harmful substances
■ Consists of a honeycomb structure of platinum, palladium and rhodium
■
Reactants
Product
Benefits of removal
Carbon monoxide
Carbon dioxide
CO is poisonous gas
Nitrogen monoxide
Nitrogen
No poisonous/ acid rain
Hydrocarbons
Carbon dioxide & water
Smog + greenhouse gas
Catalyst poison: is a substances that makes a catalyst inactive
Collision Theory & Activation Energy:
Main points:
●
●
●
Collisions must occur
Minimum Activation energy must be achieved
Right orientation required
Effective Collision: Is one that results in the formation of products
Activation Energy: Minimum energy that colliding particles must have for a reaction to occur
Reaction Profile Diagram: a graph showing the change in energy of a chemical reaction with
time as the reaction progresses
Rate of production of oxygen from hydrogen peroxide:
●
●
●
●
●
Conical flask
Graduated cylinder
Delivery Tubing
Water
Stopwatch
●
●
●
●
Water basin
Stopper
Retort stand + clamp
Beehive shelf
Reactant: Dilute Hydrogen peroxide
Catalyst: Manganese Dioxide
Product: Water
𝐻2 𝑂2 → 𝐻2O + ½ 𝑂2
Average rate:
Total volume of 𝑂2 produced / Total time for reaction to go to completion
Disappearing Cross Experiment:
Reactant: Dilute Hydrochloric acid & Sodium Thiosulfate
Product: sulphur dioxide, sodium chloride, sulphur and water
●
Swirl conical flask
Ch21: Fuels and Heats of Reaction
Organic chemistry: the study of compounds containing carbon.
Hydrocarbon: A molecule made of hydrogen and carbon atoms only.
Saturated hydrocarbon: a hydrocarbon molecule containing only single bonds.
Unsaturated hydrocarbon: a hydrocarbon molecule containing at least 1 double or triple C C
bond
Homologous Series: A series of hydrocarbons with similar chemical properties and methods of
preparation where each member differs by a CH2 group from the previous one. Each member
follows a general formula and have graduations in physical properties
Structural isomer: Molecules with the same molecular formula but a different structural
formula
Aliphatic: a straight chain or branched organic molecule (doesn’t contain a benzene ring
structure)
Aromatic molecule: an organic molecule containing a benzene ring structure
-Benzene was discovered by Micheal Faraday
-Methylbenzene also known as toluene
Properties of Benzene:
● Reactivity: Benzene is unreactive in addition reactions - tests for unsaturation does not
work
● Bond lengths: Benzene C-C bonds lengths are intermediate between that of a single bond
and double bond
● Solubility: Insoluble in water, dissolves in organic solvents
●
Other characteristics: Toxic, volatile liquid, carcinogenic
Fuels: - the Petrochemicals
Fossil fuel: Fuels formed from the remains of dead plants and animals that lived millions of years
ago.
Fractions: Group of hydrocarbons with similar boiling points
Refinery Gas ( petroleum gas)
Petrol (light gasoline)
Naphtha
Kerosene (Paraffin)
Diesel oil (gas oil)
Lubricating oil
Fuel oil
Bitumen
C1-C4
C5-C10
C7-C10
C10-C14
C14-C19
C19-C35
C30-C40
>C50
LPG in cooking
Fuel for Cars
Useful component of Plastic
Fuel for Aircraft
Fuel for Automobiles
Waxes/Polishes
Fuel for ships and industry
Material used on houses and
roads
Residue Fractions: Lubricating oil, fuel oil, bitumen
● Fractional distillation: A separation technique used for separating crude oil based on its
components’ boiling points. The column has a negative temperature gradient.
● Auto-ignition/Knocking: the premature ignition of the petrol-air mixture before normal
ignition of the mixture by a spark plug.
● The octane number : is a measure of the tendency of the fuel to resist knocking.
Reference hydrocarbons:
○ 2,2,4-trimethylpentane - iso-octane- has octane number 100
○ Heptane has an octane number of 0
Factors affecting Octane number:
○ Length of Chain:Shorter the chain, higher the octane number
○ Number of branches: Greater number of branches → higher octane number
○ Straight chain or cyclical : Cyclical and aromatic molecules have higher octane
numbers
●
●
Isomerisation: changing straight-chained hydrocarbons into branched-chained isomers
by reforming bonds
Catalytic cracking: the breaking down of long-chain hydrocarbon molecules into
●
●
short-chain molecules.
Dehydrocyclisation: Turning straight-chains into cyclical chains, producing hydrogen
molecules in the process
- E.g. Paraffin catalyst: porcelain
Adding of Oxygenates: Adding of any fuel containing oxygen - increases octane no. + is
more environmentally friendly
Main examples: ethanol, methanol, MTBE
How to produce Hydrogen gas to be used as a fuel?
●
●
●
●
Dehydrocyclisation
Steam reforming of a natural gas
Electrolysis of water
Coal Gasification.
Benefits of Hydrogen as a fuel?
● Produces high amounts of energy
●
Cleaner for the environment
Adding of Mercaptans:
Allows gas leaks to be detected
Thermochemistry:
● Heat of a reaction:ΔH the heat change that occurs when a number of moles of reactants
completely react to form its products.
● Heat of combustion: ΔHC - the heat change that occurs when one mole of a substance is
completely burnt in excess oxygen.
● The Kilogram calorific value of a fuel : is the heat energy produced when 1 kg of the fuel
is completely burned in oxygen.
Heat of combustion/Mr X 1000
● Bond energy: is the average energy required to break one mole of a particular covalent
bond and to separate the neutral atoms completely from each other.
● Heat of neutralisation: Is the heat change that occurs when one mole of a particular H+
ions from an acid reacts with one mole of OH- ions from a base.
● Heat given out = (mass) x (heat capacity) x (rise in temp)
● Standard state: this is the state that an element or compound are found in at 25C and
1atm or 101kPa
● Heat of formation: ΔHf this is the heat change that occurs when one mole of a
compound is made in its standard state from its elements in their standard states.
● Hess’s Law: 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 changes if the reaction is carried out in
one stage.
Law of conservation of energy: Energy cannot be created nor destroyed, but can only be
converted from one type into another
Heat of neutralisation:
●
React HCl + NaOH → NaCl +𝐻2𝑂
●
Heat of Reaction/number of moles = Heat of combustion
●
Calorimeter/graduated cylinder used
Standard States:
●
●
●
●
Liquid - Mercury & Bromine
Gas - Noble Gases + Fl + Cl + H + O + N
Solid- all the rest
Diatomic Elements: Hydrogen, Nitrogen, Fluorine, Oxygen, Iodine, Chlorine,, and Bromine
Preparation of Ethane
● 𝐶2𝐻4 + 𝐻2 →𝐶2𝐻6
Preparation of Ethene:
Apparatus:
●
●
●
●
Bunsen burner
2 test tubes
Tubing
Water
●
●
●
●
Water basin
Stopper
Retort stand + clamp
Glass wool
● 𝐶2𝐻5OH → 𝐶2𝐻4 + 𝐻2𝑂
●
●
●
Glass wool used to soak up Ethanol and keep it in place
𝐴𝑙2𝑂3 (a dehydrating agent/catalyst) used to slow down the vaporisation of Ethanol since if
heated directly it won’t have enough time to convert to Ethene
Move Bunsen burner occasionally towards ethene to drive vapour over catalyst
Observation:
●
Bubbles collect in test tube
Properties of Ethene:
1. Physical appearance – colourless gas with a sweetish smell.
2. Solubility: the gas is insoluble in water; Water is polar, Ethene is non-polar
Hydrogen bonds not formed
3. Combustion: Using lighted wax taper- Yellow Luminous (smoky) flame observed and
production of CO₂ - test with limewater
4. Tests for unsaturation:
a. Decolourisation of Bromine water: orange/yellow → colourless
b. Decolourisation of Dilute Potassium Permanganate (VII): Purple → colourless
Preparation of Ethyne:
Apparatus:
●
●
●
●
Flask
Tap funnel
Stopper
Water
●
●
●
●
●
Test tube/glass jar
Water basin
Retort stand
Delivery tubing
Two holed stopper and bottle
● 𝐶𝑎𝐶2 + 2𝐻2𝑂 → 𝐶𝑎2(𝑂𝐻)2 + 𝐶2𝐻2
●
●
●
Drop water few drops at a time
Stopper test tube under water
Impurities can be removed by bubbling the gas through acidified copper sulphate
Acidified Copper sulfate
Observation:
●
●
●
●
Calcium Carbide is grey black solid
Fizzing in flask and white solid formed
Bubbles collect in test tube/glass jar
Flask becomes warm
Properties of Ethyne:
1. Physical appearance – colourless gas with a sweetish smell.
2. Solubility: the gas is insoluble in water; Water is polar, Ethyne is non-polar
Hydrogen bonds not formed
3. Combustion: performed in fume cupboard - a smokier/more luminous flame is observed
+ formation of soot
4. Tests for unsaturation:
a. Decolourisation of Bromine water: orange/yellow → colourless
b. Decolourisation of Dilute Potassium Permanganate (VII): Purple → colourless
Ch17: Chemical Equilibrium
Reversible reaction: when a reaction can go in both the forward and reverse direction
Hydrocarbon: A molecule made of hydrogen and carbon atoms only.
Chemical equilibrium: A state of dynamic balance in a reversible reactions where the rate of the
forward and is the same as the rate of the reverse reaction
Dynamic equilibrium : The rate of the forward reaction is equal to the reverse reaction
Le Chatelier's Principle : If a stress is applied to a system at equilibrium, the system re-adjusts
to relieve the stress applied
Haber Process : Making ammonia (200 ATM; 500C)
The Contact Process: Making sulfuric acid (slightly over 1 atm; 450C)
Stresses:
●
●
●
Pressure:
Increase in press. → equilibrium shifts to the side with least no. of particle
Temperature:
Increase in temp. → shifts equilibrium in the direction of the endothermic reaction
Conc. of reactants:
Increase in 𝐶𝑜𝑛𝑐𝑟→ shifts equilibrium to the right
●
Conc. of products:
Increase in 𝐶𝑜𝑛𝑐𝑝→ shifts equilibrium to the left
●
Catalyst:
Addition of Catalyst → allows system to reach equilibrium at a faster rate without shifting
the system’s equilibrium
Equilibrium Constant:
●
●
Temperature dependent
Larger value of Kc the greater the extent the equilibrium is pushed towards the products
Haber Process :
●
𝑁2+ 3𝐻2 ⇌ 2𝑁𝐻3
●
Catalyst 𝐹𝑒 𝑂
2 3
- ΔH
●
●
●
●
Ammonia uses:
○ Fertilisers, explosives, cleaning agents
High Pr. / Low Temp. preferable
Too High pr. → Costly maintenance, risk of gas leaks
Too low temp. → rate of reaction is too slow
○ Compromise Temp: 500C Press: 200 atm
The contact process:
●
2𝑆𝑂2+ 𝑂2 ⇌ 2𝑆𝑂3 - ΔH
●
Catalyst 𝑉 𝑂 - Vanadium pentoxide
●
●
●
High Pr. / Low Temp. preferable
Too High pr. → Costly maintenance, risk of gas leaks
Too low temp. → rate of reaction is too slow
○ Compromise Temp: 450C Press: slightly over 1 atm
2 5
Ch18: pH and Indicators
pH = -log10 [H+]
Negative logaritithm of the conc. Of H+ to the base 10
pOH = -log10 [OH-]
Self-ionistaion of water:
+
−
𝐻2𝑂 ⇌ 𝐻 + 𝑂𝐻
+
−
𝐻2𝑂 + 𝐻2𝑂 ⇌ 𝐻3𝑂 (hydronium) + 𝑂𝐻
Ionic product of water (Kw) = [H+][OH-] (temp dependent)
pH of pure water always = 7
Acid/ Base dissociation constants
+
HA ⇌ 𝐻
+ A-
Ka = [H+][A-]
[HA]
For Weak Acids/bases:
[H+] =
√Ka x Macid
[OH-] =
√Kb x Mbase
BOH ⇌ B+ + OHKb = [B+][OH-]
[BOH]
Range of indicator: is the pH range over which there is a clear colour change
Limitations of the pH scale:
●
It does not work with concentrated solutions as once conc. goes above 1M complete
dissociation does not always occur and therefore any calculations are inaccurate
●
pH is limited to the 0-14 scale and solutions in water
Indicator:
Methyl Orange
Phenolphthalein
Litmus
pH Range:
3-5
8-10
5-8
Colours:
Red (below 3)
Colourless (below 8)
Red (below 5)
Yellow (above 5)
Pink (above 10)
Blue (above 8)
Formulae:
1 Litre= 1dm³/ 1000cm³ / 1 x 10⁻³m³
STP: 0C/ 273 Kelvin - 1x10⁵ Pascals/ 100kPa
RT:25C
(gas)1n=24dm³ at RT/ 22.4dm³ at S.T.P
(water) 1 gram= 1 ml
n= mass(g)/Ar
or
Mass(g)/Mr
n= vol(dm³) x conc (mol/dm³)
V1 x P1 / T1 = V2 x P2 / T2
T-measured in Kelvin(absolute scale)/Centigrade scale(celsius) 0C= 273K RT= 25C
P-measured in Pascals/kiloPascals
PV=nRT
(Pressure- Pa)(Vol-m³)(R-8.31 Joule mol^-1K)(Temperature-K)
The Mole: is the amount of a substance which contains 6 x 10^23 particles of that substance
Relative molecular mass :The average mass of a molecule of that compound compared with 1
twelfth of the mass of one atom of the carbon-12 isotope
Experiments:
Finding Mr of Volatile liquid:
●
●
●
●
Set up bunsen burner/tripod/ gauze- heat up beaker ⅔ full of water - keep at high temp
less than boiling
Weigh conical flask, circular piece of tinfoil big enough to cover the top of the flask and
rubber band
Put drops of x volatile liquid in flask
Seal up flask w/ tinfoil/rubber band - prick small hole in foil
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Submerge flask into the beaker of water w/ retort stand
Control the flame until liquid in flasks seems to have evaporated
Cool flask/dry measure temp of water (T2)
Weigh total mass of flask, contents/ band/ tinfoil
Find mass by subtracting recorded figures for mass of flask and co.
Find vol of flask by filling up with water and measuring it in a graduated cylinder(V2)
Use gas laws/ n= mass/mr and comparing it with STP or RT conditions
Sources of Error:
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Balance may not be accurate enough
Full flask may not be at the same temperature
Flask/foil may not be dry
Measuring cylinder may not be accurate
Finding Mr of Volatile liquid with Gas syringe:
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Air is drawn into the syringe. Self sealing cap is placed
Syringe placed in beaker of boiling water until volume of air become steady- measure
temp and initial volume
Small hypodermic syringe filled with set amount of x Volatile liquid- is weighed
Small portion of liquid injected into gas syringe self sealing cap
Hypodermic syringe reweighed and final vol of air in gas syringe noted
Subtract Use Combined gas law and n=Mass/Mr
Source of error:
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Bubbles in hypodermic syringe
Ch13: Volumetric Analysis
Standard Solution: Solution where concentration + volume is known
Primary Standard: a substance which can be obtained in a stable, pure and soluble solid form
so that it can be weighed out and dissolved in water to give a solution of accurately known
concentration
Precipitate: A solid formed out of a solution
Saturated: When no more solute can be dissolved at a certain temperature
Super Saturated: When there is more solvent than what is needed for a solution to be saturated
as a result of having been cooled from a higher temperature to a temperature below that at
which saturation occurs.
Concentrated: when there is a large amount of solute relative to the amount of solvent
Dilute: When there is a small amount of solute relative to the amount of solvent
Titrate/Analyte: Substance of unknown concentration
w/w:
-10%w/w 10g:100g
w/v:
-10%w/v
10g:100cm^3
v/v:
-10%v/v
10cm^3:100cm^3
Parts Per Million:
1mg/L = 1ppm
1000mg=1g
Molarity:
mol/dm^3
List of generic primary standards:
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NaCl (sodium chloride)
(Anhydrous) 𝑁𝑎2𝐶𝑂3 (sodium carbonate)
𝐾2𝐶𝑟2𝑂7 (Potassium chromate)
Indicators:
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Methyl Orange: Used in reactions between a strong acid and a weak base
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Phenolphthalein: Used in reactions between a weak acid and a strong base
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Acid-Red (pink)
Neutral - yellow
Basic - yellow
Acid- colourless
Basic - pink
List of Acid-Base Titrations:
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Titration of sodium carbonate against hydrochloric acid:
○ Primary Standard: (Anhydrous) 𝑁𝑎2𝐶𝑂3
○ Analyte: HCl
○ Chemical Equation:
𝑁𝑎2𝐶𝑂3 + 2HCl → 2NaCl + 𝐻2O + 𝐶𝑂2
○ Indicator: Methyl Orange
○ Colour Change: Yellow →red (Pink)
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Titration of sodium hydroxide against hydrochloric acid:
○ Standard: Standardized HCl
○ Analyte: NaOH
○ Chemical Equation:
HCl + NaOH → NaCl + 𝐻2O
○ Indicator: Methyl Orange
○ Colour Change: Yellow → red (Pink)
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Titration to determine amount of Ethanoic Acid in Vinegar:
(Vinegar is a dilute solution of Ethanoic Acid)
○ Standard: Standardized NaOH
○ Analyte: Dilute 𝐶𝐻3COOH
○ Chemical Equation:
𝐶𝐻3COOH+ NaOH → 𝐶𝐻3COONa + 𝐻2O
○ Indicator: Phenolphthalein
○ Colour Change: Pink → colourless
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Titration to determine amount of water crystallisation in washing soda:
○ Standard: Standardized HCl
○ Analyte: 𝑁𝑎2𝐶𝑂3x𝐻2O (washing soda)
○ Chemical Equation:
𝑁𝑎2𝐶𝑂3 + 2HCl → 2NaCl + 𝐻2O + 𝐶𝑂2
○ Indicator: Methyl Orange
○ Colour Change: Yellow →red (Pink)
Method:
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Weigh substance on a balance on a clock glass
Dissolve it in distilled water in beaker with glass rod
The solution is transferred into 250cm^3 volumetric flask funnel
Clock glass, beaker, glass rod, funnel washings are transferred to the volumetric flask
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Rinse the clock glass with deionised water until bottom of meniscus is brought up to the
mark
Flask is inverted many times with stopper
Wash pipette, burette and conical flask with deionised water
Titration: Lab procedure where a volume of one solution is added to another solution until
reaction is complete
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25cm^3 of standard solution is pipetted into conical flask
Burette is rinsed with a small portion of titrant
Burette attached to a retort stand is filled with the bottom of the meniscus of the HCl
(titrant) at zero mark using a funnel
Indicator added to Standard solution in the conical flask - place this flask on white tile
Slowly open Burette with thumb and two fingers
Swirl the conical flask open and close the burette at small intervals
When colour change is present close burette and record amount of HCl
Repeat
Notes:
*Copper/Chromium(Cu/Cr)- half filled sublevel
*1s,2s,2p,3s,3p,4s,3d
Cr 1s2 2s2 2p6 3s2 3p6 4s1 3d5.
Cu 1s2 2s2 2p6 3s2 3p6 4s1 3d10
*n1(or any no.) =energy levels
n= moles
* Z= Atomic No.
A= Atomic weight nuclear formula shows Z,A + chemical symbol
Cation= positive ion/ Cathode- neg electrode
Pipette
Burette
Anode=positive electrode Anion=negative ion