H2 Chemistry

Concept Outline Fu Siyang

14S7B

Sept 2015

Foundational Chemistry

One mole of a substance It is the amount of substance that contains the same number of particles as there are atoms in 12 g of 12C isotope.

Avogadro’s number It is the number of atoms in 12 g of 12C isotope.

Avogadro’s Law Equal volumes of all gases, under the same temperature and pressure, contain the same number of

particles.

Reduction A process whereby a substance gains electrons, resulting in a decrease in oxidation number

Oxidation A process whereby a substance loses electrons, resulting in an increase in oxidation number

Redox reaction A reaction that involves reduction and oxidation simultaneously

Disproportionation A redox reaction in which the same substance is both oxidised and reduced

Oxidation number The number of electrons to be added or subtracted from an atom in a combined state to convert it to

elemental form

Atomic orbital The region of space with a 90% probability (or more) of finding an electron

Effective nuclear charge The net nuclear charge experienced by an outer electron

Chemical Bbonding

Metallic bonding The electrostatic forces of attraction between metal cations and the sea of delocalised electrons in a

metal

Ionic bonding The electrostatic forces of attraction between oppositely charged ions in an ionic compound

Covalent bonding The electrostatic forces of attraction between the nuclei of atoms and their shared pair of electrons

The VSEPR theory 1) The electron pairs around a central atom will adopt a geometry that would minimise the

repulsion between them by placing themselves as far away as possible in a 3-dimensional space;

2) Lone pair-lone pair repulsion > lone pair-bond pair repulsion > bond pair-bond pair repulsion

Non-polar molecule One that has no overall dipole after resolving the dipole moments of every covalent bond in the

structure

Polar molecule One that has a net dipole moment

Chemical Energetics and Thermodynamics

Activation Energy The activation energy of a reaction is the minimum energy which the reacting particles must possess in

order to overcome the energy barrier before becoming products

Standard enthalpy change of reaction The enthalpy change when molar quantities of reactants as specified by the chemical equation react to

form products at 1 atm and 298K.

Standard enthalpy change of formation The enthalpy change when 1 mol of a substance is formed from its constituent elements in their

standard states at 298K and 1 atm.

Standard enthalpy change of combustion (-ve) The heat evolved when 1 mol of a substance is completely burnt in excess oxygen at 298K and 1 atm.

Standard enthalpy change of neutralisation (-ve) The heat evolved when 1 mol of water is formed in the neutralisation reaction between an acid and a

base, at 298K and 1 atm.

Standard enthalpy change of atomisation for an element (+ve) The energy required when 1 mol of gaseous atoms is formed from the element at 298K and 1 atm.

Standard enthalpy change of atomisation for a compound (+ve) The energy required to convert 1 mol of the compound into gaseous atoms at 298K and 1 atm.

Standard enthalpy change of hydration (-ve) The heat evolved when 1 mol of free gaseous ions is dissolved in an infinite volume of water at 298K and

1 atm.

Standard enthalpy change of solution The enthalpy change when 1 mol of solute is completely dissolved in an infinite volume of solvent at

298K and 1 atm.

Bond dissociation energy (+ve) The energy required to break 1 mol of a particular covalent bond in a specified molecule in the gaseous

state.

Bond energy (+ve) The average energy required to break 1 mol of a covalent bond in the gaseous state.

Lattice energy (-ve) The heat evolved when 1 mol of solid ionic compound is formed from its constituent gaseous ions.

The first ionisation energy (+ve) The energy required to remove 1 mol of electrons from 1 mol of gaseous atoms to form 1 mol of singly

charged gaseous cations.

The first electron affinity (1st –ve; 2nd +ve) The enthalpy change when 1 mol of electrons is added to 1 mol of gaseous atoms to form 1 mol of singly

charged gaseous anions.

The Hess’s Law The enthalpy change of a reaction is determined only by the initial and final states and is independent of

the reaction pathway taken.

Entropy It is a measure of the randomness or disorder in a system, reflected in the number of ways that

molecules and energy in a system can be arranged.

Enthalpy Change The difference in energy level between the products and reactants

Spontaneous process A spontaneous process is one that, in the absence of any barrier such as activation energy, takes place

naturally in the direction stated.

The Gaseous State

Ideal gas It is a hypothetical gas of which pressure-volume-temperature behaviour can be completely accounted

for by the ideal gas equation pV=nRT.

Partial pressure Partial pressure of a gas is the pressure exerted by that individual gas in a mixture of gases on the sides

of a container.

Dalton’s Law of Partial Pressures The total pressure of a mixture of gases is the sum of the partial pressures of the constituent gases.

The kinetic theory of gases (important assumptions) 1) The gas particles have negligible volume compared to the volume of their container

2) The intermolecular forces between gas particles are negligible

3) Collisions between gas particles and collisions between gas particles and walls of the container

are perfectly elastic

Chemical equilibria

Le Chatelier’s Principle When a system in equilibrium is subjected to a change on conditions which disturbs the equilibrium, the

position of equilibrium will shift in a way so as to reduce that change.

Dynamic equilibrium When the rate of the forward reaction equals the rate of the reverse reaction, the system is in dynamic

equilibrium.

The equilibrium law If a reversible reaction is allowed to reach equilibrium, the product of the concentrations of each

product (raised to the appropriate powers) divided by the product of the concentrations of each

reactant (raised to the appropriate powers) has a constant value called the equilibrium constant, K, at a

constant temperature.

𝐾𝑐 =[𝐶]𝑐[𝐷]𝑑

[𝐴]𝑎[𝐵]𝑏

Reaction Quotient At any instant,

𝑄𝑐 =[𝐶]𝑐[𝐷]𝑑

[𝐴]𝑎[𝐵]𝑏

The Haber Process 1) 450 ℃

2) 250 atm

3) Finely divided Fe catalyst

Ionic Equilibria

Bronsted-Lowry acid-base theory 1) An acid is a proton donor.

2) A base is a proton acceptor.

Lewis acid-base theory 1) An acid is an electron pair acceptor.

2) A base is an electron pair donor.

Conjugate acid-base pair A conjugate acid-base pair constitutes two species which differ from each other by one proton.

Degree of dissociation The fraction of molecules which is ionised into ions in water is called the degree of dissociation.

Salt hydrolysis A reaction in which ions react with water to produce OH- or H3O+ ions.

Buffer solution A solution that is able to resist pH changes upon addition of small amount of acid or base.

1) Acidic buffer: an aqueous solution of a weak acid and its salt

2) Alkaline buffer: an aqueous solution of a weak base and its salt

Choice of indicators Working range of an indicator is pH=pKIn ± 1.

Indicator Working range Before equivalence point

At equivalence point

After equivalence point

Methyl orange 3 – 5 Red Orange Yellow

Phenolphthalein 8 – 10 Colourless Pink Red

Thymol Blue 8 – 10 Yellow Green-grey Blue

Solubility Product

Solubility of a salt The amount (or mass) of solute (salt) that can be dissolved in 1 dm3 of a given solvent to form a

saturated solution at a given temperature.

Supersaturated solution A solution that contains more solute than the amount needed to form a saturated solution.

Solubility Product Ksp of a sparingly soluble salt is the product of the molar concentrations of the constituent ions in a

saturated solution, raised to the appropriate powers at a given temperature.

Ionic product IP of a sparingly soluble salt is the product of the molar concentrations of the constituent ions in the

solution, raised to the appropriate powers at a given temperature.

Electrochemistry

Standard electrode potential The standard electrode potential of a half-cell is the potential associated with a given half-reaction when

all components are in their standard states, measured relative to the standard hydrogen electrode.

Standard cell potential The potential difference between the two half-cells measured at standard conditions.

Electrolysis A process that involves the lysing or splitting of a substance, often to its component elements, by

supplying electrical energy. The electrical energy from an external source causes a non-spontaneous

redox reaction to occur.

Electrolyte An electrolyte is a liquid or an aqueous solution, which contains mobile ions and therefore can conduct

electricity.

Reaction Kinetics

Transition State The arrangement of atomic nuclei and bonding electrons at the maximum potential energy.

Rate of reaction The change in concentration of a reactant or a product per unit time.

Half life The time taken for the concentration of a reactant to decrease to half of its initial value.

Clock reaction A clock reaction is produces a sudden change (usually visual) when a small fixed amount of a product is

formed.

Catalyst A catalyst is a substance that speeds up a chemical reaction by providing an alternative reaction

pathway with lower activation energy, while remaining chemically unchanged at the end of the reaction.

Autocatalysis One of the products of the reaction is a catalyst for the reaction itself.

Enzymes Enzymes are biological catalysts that speed up reactions in living systems while remaining unchanged at

the end of the reaction. They are specific in their action.

Substrate The reactant that binds to the enzyme’s active site is the substrate.

Chemical Periodicity

Flame Colour (in oxygen) Na Brilliant yellow flame

Mg Bright white light

Al NA (oxidised to Al2O3)

Si NA (oxidised to SiO2)

White phosphorous P4 (trigonal pyramidal structure) Bright white flame

Red Phosphorous P Soft orange flame

S Pale blue flame

Cl NA

Ar NA

Chloride salts NaCl White ionic solid

MgCl2 White ionic solid

AlCl3 Pale yellow powder that fumes in moist air; sublimes to form Al2Cl6

SiCl4 Colourless volatile simple covalent liquid

PCl3 Colourless fuming simple covalent liquid

PCl5 Off-white solid

Colours of halogens Element In gaseous

phase In liquid phase

In solid phase

In aqueous phase In non-polar solvent phase

F2 Pale yellow NA NA NA Very pale yellow

Cl2 Yellow-green

NA NA Colourless (dilute) Pale yellow (conc.)

Very pale green

Br2 Reddish brown

Reddish brown

NA Yellow (dilute) Orange (conc.)

Orange to reddish brown, depending on concentration

I2 Dark purple NA Shiny black

Pale yellow (in water) Brown (in I- (aq))

Violet

Organic chemistry

Electrophile An electron deficient species that accepts an electron pair from an electron rich species in a reaction to

form a new covalent bond

Nucleophile A species that donates an electron pair to an electron deficient species in a reaction to form a new

covalent bond

Intermediate A definite species that exists for a finite length of time in a reaction.

Free radical An atom or group of atoms that has an unpaired electron

Carbocation A species that contains a carbon atom bearing a positive charge

Carbanion A species that contains a carbon atom bearing a negative charge

Inductive effect The withdrawal or donation of electrons through a σ bond due to electronegativity and the polarity of

bonds in functional groups

Resonance effect The withdrawal or donation of electrons through a π bond due to overlap of a p orbital on the

substituent with a p orbital on the adjacent double bond or aromatic ring

Enantiomer Non-superimposable mirror images

Geometric isomerism in alkenes Geometric isomerism arises in alkenes due to restricted rotation about the C=C bond.

Proteins/Amino acids

Zwitterion Dipolar ions with no overall electrical charge

Isoelectric point The pH at which the overall net charge on the amino acid is 0 and it exists primarily as the neutral

zwitterion

Electrophoresis Separation of different amino acids at a certain pH level using an electric field

Transport proteins Move metabolites around the cell or around the whole organism

Regulatory proteins (hormone) Coordinate the activities in the body by controlling the level and type of cell functions including

metabolism and reproduction

Defensive proteins (antibody) Produced by specialised cells of the immune system in response to foreign antigens

Catalytic proteins (enzyme) Accelerate metabolic processes which produce energy, build up new cell structures and destroy old ones

Primary structure It is the sequence of amino acid residues in a polypeptide chain through formation of peptide linkages.

Secondary structure It is the regular coils and folds in localised segments of a polypeptide chain, stabilised by hydrogen

bonds formed between the lone pair electrons on the oxygen atom of a C=O bond in a peptide linkage

and the H atom in a N-H bond in another peptide linkage on the polypeptide backbone.

Tertiary structure It is the 3-dimensional shape of a polypeptide chain, stabilised by side-chain interactions.

Quaternary structure It is the overall special arrangement of several polypeptide chains through side chain interactions.

Denaturation It is the process of destroying the secondary, tertiary and quaternary structures and altering the specific

conformations of a protein, causing it to lose its ability to perform its function.

Transition elements

Transition element It is a d-block element that forms some compounds containing its ion with an incomplete d-subshell

Complex A complex is formed by a central metal atom or cation dative-bonded by surrounding anions or

molecules (known as ligands)

Ligand A ligand is an ion or molecule that has at least one lone pair of electrons that can be donated into the

vacant orbitals in the central metal atom or ion

Coordination number The coordination number indicates the number of dative bonds around the central atom or ion.

Explanation Questions

D-d electronic transitions In a free metal atom or ion with partially filled 3d orbitals, the 5 3d orbitals are degenerate. In the

presence of ligands in an octahedral field, the 5 3d orbitals split into two different energy levels with an

energy gap ΔE in between. There are vacancies in the higher energy d orbitals, and the promotion of an

electron from a lower energy level to a higher one requires the absorption of a photon with energy ΔE,

usually in the visible region of the electromagnetic spectrum. Such d-d transition is responsible for the

colour of the complex ion, and the colour observed is the complement colour of the photons absorbed.

Toxicity of CO Haemoglobin is the Fe-containing protein in red blood cells that transports oxygen in blood. A

haemoglobin molecule in the lungs picks up an O2 molecule and forms reversible dative binding between

the Fe atom and the O2 molecule, enabling haemoglobin to carry oxygen around the body and release it

when needed. However, CO undergoes ligand exchange with oxygen and binds with the Fe atom in

haemoglobin irreversibly to form a very stable complex carboxyhaemoglobin. The affinity of human

haemoglobin for CO is much larger than that for O2, thus a small quantity of CO can inactivate a large

number of haemoglobin molecules for oxygen transport. If the concentration of CO is high and the level

of carboxyhaemoglobin is too high, oxygen transport is shut down and death occurs.

Thermal Stability of Group II compounds The thermal stability of Group II nitrates, carbonates and hydroxides increases down the group. This is

because the cationic radii increase down the group, while charges remain the same. Charge density of

the cations decreases down the group, thus polarising power of the cations decreases down the group.

They are less able to distort the electron cloud of the anion, weakening the covalent bonds within the

anion to a smaller extent. Covalent bonds within the anion are less likely to be broken down the group,

hence higher temperature (higher energy) is required to break the bonds.

Abnormally low bond energy of F2 From Cl to I, the atomic radii increase and hence the orbital overlap between the atoms becomes less

effective. The bond length increases and the X-X bond becomes weaker.

The F-F bond is so short that the lone pairs of electrons on the F atoms repel each other and weaken the

F-F bond.

Abnormally small (in magnitude) first electron affinity of F atom From Cl to I, the number of quantum shells increases. Any electron added to the atom is less strongly

attracted to the nucleus. The ability of the halogen atom to accept the electron decreases, thus first

electron affinity becomes less exothermic.

F is an exception due to its extremely small size of the atom. The electron is placed into a region of space

that is crowded with electrons, causing strong repulsion. Hence the electron added is less strongly

attracted to the nucleus.

Electronic configurations of Cr and Cu The electronic configuration of Cr is [Ar] 3d5 4s1 because equal distribution of electronic density in the 3d

subshell results in the overall reduction of inter-electronic repulsion in the atom.

The electronic configuration of Cu is [Ar] 3d10 4s1 because symmetrical electron distribution in the 3d

subshell is preferred.

Surface lustre of metals The delocalised electrons do not reside in any orbitals. When light falls on the surface, most photons

simply bounce off as the metal is unable to absorb light. Most of the light is reflected and observed,

accounting for the shininess of the metal.

Deviation of actual gases from ideal gas behaviour At low temperature, intermolecular forces of attraction becomes significant and collisions become

inelastic. When the particles do not have enough energy to overcome intermolecular forces of

attraction, liquefaction occurs.

At high pressure, intermolecular distances become shorter and intermolecular forces become

significant. The impact of molecules striking on the walls is lowered and the pressure is lower than what

is expected of an ideal gas. The molecules also take up a large portion of the volume of the container,

hence the assumption that volume of the molecules are negligible compared to the volume of the

container is no longer valid.

Mode of action of heterogeneous catalysts 1) Adsorption: the reactant molecules diffuse towards the surface of the solid catalyst and are

adsorbed onto the active sites at the surface. The adsorption process brings reactant molecules

closer together, increasing their surface concentrations and collision frequency. The process also

weakens the bonds in the reactant molecules, thus lowering the activation energy. The reactant

molecules are also correctly oriented for reaction to occur.

2) Desorption: product molecule desorbs and diffuses away from the catalyst surface. The active

sites become free to adsorb new reactant molecules. When the active sites are fully occupied

(i.e. saturated), the order of reaction with respect to the reactant molecules becomes 0.

Relationship between [substrate] and rate of an enzyme-catalysed reaction At low [substrate], the enzyme concentration is greater than substrate concentration, so as more

substrate is added, the rate of reaction increases. The reaction is first order with respect to the

substrate.

At high [substrate], when all the active sites are filled with substrate molecules (i.e. saturation point),

any further increase in [substrate] does not affect the rate. The reaction is now 0th order with respect to

the substrate.

Mode of action of a buffer (CH3CO2H and CH3CO2- Na+)

The solution contains large reservoirs of unionised CH3CO2H and CH3CO2- ions from CH3CO2

- Na+.

When a small amount of H+ is added to this buffer, the H+ ions react with the large reservoir of the

CH3CO2- ions and are removed and pH remains almost constant.

When a small amount of OH- is added to this buffer, the OH- ions react with the large reservoir of the

CH3CO2H molecules and are removed and pH remains almost constant.

The importance of a salt bridge in the electrochemical cell Without a salt bridge, charge imbalance is given rise in the cell and current will cease. In the salt bridge,

anions in the salt bridge are released into the anode solution to balance the excess cations, and some

cations from the solution will also enter the salt bridge; cations in the salt bridge are released into the

cathode solution to balance the excess anions, and some anions from the solution will also enter the salt

bridge. The salt bridge is necessary to complete the circuit and to maintain electrical neutrality.

Electrolysis of concentrated NaCl solution For concentrated NaCl solution, the higher concentration of Cl- causes the position of equilibrium

Cl2 + 2 e- <-> 2Cl-

to shift to the left according to Le Chatelier’s Principle. This results in E(Cl2/Cl-) becoming less positive

than EΘ(O2/H2O). Thus Cl- is preferentially oxidised at the anode instead of H2O.

At the cathode, the difference between EΘ(Na+/Na) -2.71 V and EΘ(H2O/H2) -0.83 V is too large; even

concentrated solution cannot make the value of EΘ(Na+/Na) to be more positive than -0.83 V. Water is

still preferentially reduced.

Electrolytic purification of copper (half equations are omitted) EΘ(Ni2+/Ni) is the least positive of the four values, and EΘ(Cu2+/Cu) is the 2nd least positive. Ni and Cu are

preferentially oxidised at anode. The anode dissolves. Ag in the impure copper is not oxidised due to

unfavourable electrode potentials, hence it sinks to the bottom once Cu dissolves as the “anode sludge”

and is collected.

At the cathode, since EΘ(Cu2+/Cu) is the most positive, Cu2+ will be preferentially reduced and Cu metal

will deposit on the cathode. Ni2+ and H2O will not be reduced due to unfavourable electrode potentials.

Geometric isomerism affecting melting/boiling points The cis isomers tend to have higher boiling points because they tend to be more polar.

The trans isomers tend to have high melting points because they pack better due to their high

symmetry. The dispersion force work more effectively in holding the molecules together, hence more

energy is needed to overcome the forces of attraction.

Effects of CFCs on the ozone layer In the stratosphere, oxygen atoms are produced when O2 atoms absorb UV light at 250nm wavelength.

Once oxygen atoms form, they can react with oxygen molecules to produce ozone, which by absorption

of UV light, decomposes to reform oxygen atoms and molecules. An equilibrium is set up. CFCs act as a

homogeneous catalyst for the destruction of ozone, and the process uses up the oxygen atoms needed

to make more ozone by the natural ozone formation reaction.