IB chemistry revision notes

Electronegativity in the Periodic Table

• The Alkali Metals (Group I)

• These elements have low values for electronegativity
• As you go down the group, electronegativity decreases, even thought the number of protons is increasing.

• The Halogens (Group VII)

• These elements have relatively high electronegativity values.
• As you go down the group, electronegativity decreases, even thought the number of protons is increasing.

• Across a Period

• Electronegativity increases across the Periodic Table from left to right.
• This is due to the fact that the nuclear charge increases and the atomic radius decreases slightly.

• Down a Group

• Electronegativity decreases because although Z increases, the effect is more than compensated by the increase in screening electrons and the diluting effect of atomic radii getting larger.
• Another way of saying this is that the outer shell is further away from the nucleus and so the shielding increases.

• Noble Gases

• These elements have no desire to gain electrons; therefore they do not have electronegativity values.
• Argon and Xenon can form a bond. So these as the bottom of the group could have electronegativity values, but those at the top can’t.

• Halogens and Halides

• Halides are the ions of the Halogens
• The more electronegative the halogen elements are, the more able they are to pull the electron off the halide ions which are lower in the group than themselves.

E.g.

• The chlorine molecule is reduced- it is the oxidising agent. The bromide ions are oxidised, they’re the reducing agents.

Ionisation Energy

• First Ionization Energy→ the energy required to remove one electron from each atom of a mole of atoms in the gas state, to form one mole of cations in the gas phase, under s.t.p
•  is positive.

Across a Period

• In every period the noble gas has the highest value.
• Going across the Periodic Table, the general trend is a rise in. This is because the outer electrons are going into the same electron shell, but the nuclear charge in increasing. This means they are held more tightly and thus, more energy is required to remove them.
• Although the general trend is up, elements in the 2nd and 5th groups have a higher than normal value for because the ones in Group II have a full s shell while those in Group V have a half full p shell. This means that we see a drop in the 3rd and 6th elements.

Down a Group

• The ionisation energy decreases as it is easier to pull out the electron because:

• The outer electron is further away from the nucleus. This outweighs the increase in Z, thus the electrostatic forces fall (since they are inversely proportional to the square of the distance).
• The number of electrons in the inner shells of the atom increases, thus increasing shielding.

 These two factors make the energy required to pull one electron off smaller.

Successive Ionization Energies

• These are the ionisation energies required to remove more than one electron from a single atom.
• Electrons are removed from the outside outwards. Special rules exist for the d shell.

1. The energy increases as each electron is removed from the same shell, for the ion has become more positive (and smaller, so there are 2 reasons why the electrostatic forces get larger).
2. There is a large increase in the energy as a new shell of electrons is broken. This is because electrons are closer to the nucleus and there are fewer electrons shielding electrons.
3. The step increases on the energy change increases as we get closer to the nucleus.

Tips for drawing:

• The electronic structure backwards
• Each “flat” set gets slightly steeper each time.
• The “steps” get slightly bigger each time.

Melting Points

• The m.p.t depends on the type of bonding. The bonding depends on the arrangement of the outer electrons.
• Going across a period shows a behaviour going from metallic to non metallic.
• Having one, two or three electrons in the outer shell encourages metallic behaviour. When we get to the half filled shell we get a giant covalent structure. The next elements are gasses, being made of simple covalent molecules. The eighth element forms a monoatomic gas.

• The Alkali Metals (Group I)

• As we go down the group the m.p.t decreases. Also, metals are softer.
• All the elements have a body centred cubic structure made up of cations and delocalised electrons. As we go down, the radius of the ions increases and so the cations are further away from the delocalised electrons, therefore the attraction is weaker.

• The Halogens (Group VII)

• As we go down the group, the m.p.t increases.
• The effect is due to increased Van der Waal’s forces between the molecules- the higher the number of electrons, the higher the possible forces between the temporary dipoles.

Chemical Properties

• The Alkali Metals (Group I)

• All are soft and have low densities.
• All have coloured flame tests.
• They’re very electropositive: like to make cations easily.
• Behave as very strong reducing agents, thus oxidising themselves.
• Metal (I) hydroxides are very soluble. E.g. KOH
• Metal (I) hydroxides are very strong bases. E.g.
• Reactivity increases down the group, due to the easiness with which the electron can be lost from the metal. (Ionization energy decreases).
• Reaction with water:

Removing the Spectator Ions:

• Reaction with Halogens:

• All Group I metals react with halogens to form alkali metal halides.
• Alkali metal + Halogen → Alkali metal Halides

• The Halogens (Group VII)

• They are the salt makers.
• As we go down the group, atomic, ionic and covalent radii increase. Ionisation energies for this group are the second highest in the periodic table.
• As we go down the group, the reactivity of the halogen decreases, but the halide’s increases.
• As we go up the group        →halogen molecule becomes a stronger oxidising agent.

→halide ion becomes a weaker reducing agent.

• The further up the group the halogen, the better the oxidising agent. The better down, the better the reducing agent.

Use of Standard Electrode Potentials

• Electrode potentials are a measure of how strong oxidising or reducing an agent is. They predict whether a redox reaction is energetically favourable. The hydrogen half-cell was used as the base for all the S.E.P.s

• Place the most negative value at the top
• Place the most positive value at the bottom.

Test for Halide Ions

1. Add nitric acid. This will prevent other pseudo halides from reacting with the silver ions in the silver nitrate, forming a white p.p.t. E.g., which is a white p.p.t.
2. If there is no noticeable change, add silver nitrate. If chlorine ions are present, and instant whit p.p.t. of silver chloride will be produced.

In the presence of light, the white p.p.t will decompose and turn grey.

The silver chloride p.p.t. dissolves easily in ammonium hydroxide (ammonia).

3. In the case of Bromide Ions, a cream p.p.t. is produced which also decomposes with light and turns grey. However, it doesn’t dissolve in ammonia.
4. In the case of Iodide ions, a yellow p.p.t is produced which does NOT decompose with light.

Trends across the Third Period

Bonding

Ionic Bonding

• When are two elements likely to combine to form a Binary Compound which is ionic in nature?

• An ionic bond is formed by the transfer of electrons to form 2 oppositely charged particles (ions) which are held together by electrostatic forces.

Covalent Bonding

• Generally between 2 non-metals that need to gain electrons to fill their outer shells. This leads to the valence orbitals overlapping so that 2 atoms can share a pair of electrons.
• A covalent bond is the electrostatic attraction between the nuclei of two atoms and a pair of shared electrons.
• A covalent bond is the rough approximation to a spring. The longer away the more attraction; if made to go too close they repel.

• Lewis structures

• Lewis states the idea that atoms tend to bond in order to have eight electrons in the outer shell. This idea became known as the octet rule.

E.g. Fluorine:     or    

• Covalent Bond Lengths

• The minimum distance between the two nuclei of the atoms is 1 covalent bond length.

•  The bond length gets smaller→

• E.g. Ethanoic Acid

        C=O    1.122 nm                        C-O     1.134 nm

• As you go down the group, the covalent bond length increases.
• As you have more bonds, the bond dissociation energy increases, since bond strength increases.

• Electronegativity

• If two atoms have similar electronegativity then they will form covalent bonds.
• If they differ by more than 1 unit, they are a polar covalent bond.
• If they differ by more than 2 units, the bond is ionic.

• Properties

• Outside the molecule → intermolecular forces are weak.
• Inside the molecule →intermolecular forces are strong.
• Low m.p.t. and b.p.t.
• At room temp → often gas or liquid.
• Do not conduct electricity.
• Dissolve in organic solvents.
• Soft if solids.

• Polar Covalent Bonds

• Electrons in a covalent bond aren’t shared equally.
• Atoms with higher electronegativities have a greater control over electrons. This leads to a non symmetrical electron distribution.

2.1         H                                Br        3.0

V.S.E.P.R. Theory

• Valence Shell Electron Pair Repulsion Theory

• Attempts to explain the shapes of simple molecules or simple ions.
• Basic Idea→ molecules or ions have the shape which is most stable.
• To be stable pairs of electrons like to arrange themselves to be as far away as possible from other pairs.

• Rules for working out the structure:

1. Identify the central atom.
2. Work out how many electrons there are in the central atom.\
3. Count the number of atoms bonded to the central atom ad then add that number to the number of valence electrons.
4. If a species has a positive charge take away 1. If it is negative add one.
5. Divide total electrons by 2 → gives the number of electron pairs.
6. Count the number of bonds and with the number of electron pairs calculate how many are bond pairs and how many are lone pairs.

• Shapes:

• AB, AB2 → Central atom (A) has no lone pairs and one/two bonding pairs. It is a linear molecule, for 180° is the furthest possible position two atoms can be. Thus, thus are at their safest minimum repulsion and maximum energetic stability. Eg. CO2.

• AB3→No lone pairs and three bonding pairs. The shape is Trigonal planar, since all 4 atoms are on the same plane. The bond pairs are polar but the molecule isn’t because it is symmetrical. The bond angle is 120°. E.g. Boron Trichloride.

• AB4 →No lone pairs and four bonding pairs. This is a tetrahedral shape. Bonds are slightly polar but the molecule isn’t (symmetrical). The bond angle is 109° 27’ or 109.5°. E.g. Methane.
• ÄB2 →This molecule has one lone pair and two bonding pairs. The lone pair of electrons repel other bonds further away than normal bonds. This gives a V-shaped molecule that is polar. The bond angles between the lone pair and a bond is >120°. The bond angle between the two bonds is <120°. E.g. sulphur dioxide.
• ÄB3 →This molecule has one lone pair and three bond pairs. The lone pair creates a maximum repulsion. The shape it trigonal pyramidal. The bonds are polar, and because of the orientation of the bonds, the molecule is also polar. The bond angle is of ~107°. E.g. Ammonia.
• :ÄB2 → This molecule has two bonding pairs and two lone pairs. The two lone pairs will give a greater repulsion. We obtain a V-shaped molecule. The bond angle is of ~105°. E.g. Water.
• Ethane, C2H6 → Each end is a tetrahedral with bond angles of 109.5°.
• Ethene, C2H4 → Each of the H and H are less than 120° due to the repulsion of the double bond. The molecule is flat (planar) due to the rigidness of the double bond.

Intermolecular forces

• All the forces are electrostatic:

• Van der Waal’s forces:

• Electrons in atoms, molecules and ions are free to move around within the atoms.
• Van der Waal’s forces are caused by the random, instantaneous movements of electrons in a species. When this happens it can become a temporary dipole; if another species happens to be fairly close then electrons in it will be induced to move away. A temporary dipole then induces another temporary dipole. These weak forces are the Van der Waal’s forces

                   *        

            H  *  ||||||      H                →                H  *            H  *

• Directly proportional to the number of electrons in the species
• The Halogens: as we go down Group VII the atoms have more electrons. So, they can have a greater charge imbalance and induce greater dipoles. The attraction is greater and requires more heat to break.
• The Alkanes: as the number of carbon atoms increases, there are more places where temporary induced dipoles can happen. Thus, Van der Waal’s forces are stronger. More energy is required to separate atoms.

   C6H14

C4H10

    CH4

• Dipole-Dipole Attraction        

• Some molecules, due to the electronegativities of their atoms and/or their asymmetric shape have permanent dipoles.
• The dipoles attract each other. Negative ends line up with positive ends.
• Molecules that have dipole-dipole interactions have higher m.p.t and b.p.t. than similar molecules with similar masses that don’t have permanent dipoles. E.g. 1,4 dichlorobenzene → 1,2 dichlorobenzene
• E.g. Ethanal,.

    _ _ _ _ _ _ _    

• Hydrogen Bonding

• Occurs when hydrogen in bonded directly to a small, highly electronegative element (F, O, N). Its single electron is pulled closer to the other atom, leaving the H nucleus exposed (no screen). This is an extreme case of dipole-dipole bonding.
• Hydrogen bonding is the attraction between the lone pairs of electrons of a very electronegative atom and the exposed nucleus of another hydrogen compound.

    |||||                            |||||

• E.g. Water, Ammonia.        

                   

Metallic Bonding

• The bonding is the attraction between cation and delocalised electrons.

• Down Groups I and II (Alkali and Alkali Earth Metals)

• The boiling points decrease as the metallic bond becomes weaker.
• Each element donates electrons to form the sea of delocalised electrons. However, as the number of shells increases, the distance between the nucleus and the electron sea increases. Thus, the attractive force decreases and less energy is required to break the bonds.

• Across Na→Mg→Al

• The m.p.t increases as there are more valence electrons. This gives a greater difference in charge between the cations and the sea of electrons.
• Also, the ionic radii decreases, reducing the distance between the cations and the sea of electrons.

• Properties of Metals

• Good conductors of electricity → delocalised electrons.
• Good conductors of heat→electrons jumping through cations and moving energy from to .
• Shiny → light absorbed by electrons and re-emitted at different Energy levels.
• Malleable → pushing layers.
• Ductile → moving the layers.
• Grey coloured → release all ; white light intensity is low making the metal look grey.

Solubility

• In water:

• We need covalent molecules which can interact with the water molecules easily.
• Need to be the same types of intermolecular forces. E.g. hydrogen bonding or, to a lesser extent, dipole-dipole attraction.

• In organic solvents:

• E.g. Benzene
• Are generally non-polar, therefore are used to dissolve non-polar molecules that have Van der Waal’s.

Transition Metals

• Ti        V        Cr        Mn        Fe        Co        Ni        Cu        Zn
• They can form different ions.

• Properties

• Non-Volatile
• Solid at room temperature (except Mercury)
• Crystalline solids due to regular arrangement of the ions. E.g. NaCl → forms a 3D lattice where each is surrounded by six, and each  is surrounded by six. It is a face centred cubic arrangement with 6:6 coordination.
• Conducts electricity when molten or dissolved in aqueous solution. Here, ions are free to move around and transfer charge.
• High m.p.t and b.p.t→ since ions are held together in crystalline lattice by very strong ionic bonds a lot of Energy is required to separate them. This depends on three factors:

1. Size of ions: larger ions have a lower m.p.t as the coulomb force between them is lower:
2. Charge of ions: the higher the charge the greater the attraction between them and thus they require more energy to separate.
3. Type of lattice structure: packing ions determines how close, on average, the ions are to each other.

• Brittle→ when a force is applied the lattice varies its arrangement, provoking repulsion between the ions.
• Soluble in polar solvents; insoluble in non polar solvents.

• Usually requires differences of electronegativities larger than ~2.

States of Matter

Solids

• Fixed shape- particles vibrate always (more at high temperatures)
• Fixed Volume
• Particles packed close together- hard to compress
• Strong forces between the particles
• Low Entropy

Liquids

• Particles can exist in clusters and swap positions- flows, takes container’s shape.
• Fixed volume.
• Particles still touching- difficult to compress.
• Weaker forces between the particles- expends on heating.
• Higher entropy than solids

Gases

• Particles move at random (Brownian motion) - no fixed shape.
• Forces are very small- expand a lot on heating.
• Fills the space available.
• Easy to compress.
• High entropy.

Endothermic Processes

• Melting

• Particles are absorbing energy ()
• At the m.p.t. the temperature of a pure substance doesn’t increase.

• Boiling

• We are pulling the particles apart (W= Fd). As we have more space, the external energy is greater.
• At the b.p.t the temperature of a pure substance doesn’t increase.

• Sublimation

• From a solid to a gas.
• A lot of energy is required.

Exothermic Processes

• Condensation

• Particles are moving together.
• The structure becomes more organised.

• Freezing

• Releases energy to the environment.

Diffusion

• The random movement of particles from areas of high concentration to areas of low concentration.
• Tends to be faster in gases, slower in liquids and very slow in solids.

Kinetic Theory

• Pressure → The measure of the number and speed of the molecules hitting the walls of the container per second.

• If the molecules hit the container walls more often then
• If the molecules hit the container walls with greater speed then

• Temperature → a measure of the kinetic energy of the particles (the speed of movement)

• Kelvin → Absolute Temperature.
• Absolute zero is the coldest temperature→ particles stop moving completely.

• The Ideal Gas Equation.                

• n → number of moles.
• R → ideal gas constant
• T → Temp in Kelvin

• Boyle’s Law

• Temperature constant
• For a fixed mass of gas the V is inversely proportional to P.
• P1V1 = P2V2

• Charles’s Law

• Pressure constant
• For  fixed mass of gas the V is directly proportional to T

• Pressure Law

• Volume Constant

Maxwell-Boltzmann Energy Distribution Curves

1. Raising temperature shifts the peak to the right (not as tall, but wider)
2. Number of molecules with the most “popular” energy is smaller.
3. Area under both curves is the same.

Energetics

Energy Profile of Reaction

• Exothermic reactions → release heat energy to the surroundings.

1. Activation Energy EA

• The products have less energy in their bonds than the reactants, meaning that they have lost it to their surroundings.

• Endothermic Reactions → gains heat from the surroundings.

1. Activation Energy  EA
2. is positive

• The products have more energy in their bonds than the reactants, meaning that they have gained it from their surroundings.

• In all reactions we first have to put energy in to break the bonds. Then, energy is released as new bonds are made.
• If energy is lost (exothermic reaction), then the products are more stable than the reactants. In endothermic reactions, the reactants are more stable then the products
• Breaking bonds →endothermic, +
• Making bonds → exothermic, -

Hess’ Law

It states that if a reaction can take place by more than one route, the overall enthalpy change is the same whichever route is followed.

Enthalpies of Reactions

• This is the energy change (heat) at constant pressure for any stoichiometric equation.
• Using bond enthalpies:

• The bond enthalpy is the average energy needed to separate the two atoms in a bond in the gas phase.
• Given as average because the bond strength varies slightly depending upon the surrounding molecule.

• These calculations can be unreliable as the products and reactants have to be in the gas phase. They don’t take into account any energy changes associated with changes of state in either products or reactants.

• Using Enthalpies of Formation

• Enthalpy of formation is the heat energy required to make 1 mole of a substance at S.T.P. from it’s elements at S.T.P.
• S.T.P. → 25° and 1 atm.
• Ethane, C2H6 →

• Using Enthalpies of Combustions

• Enthalpy of combustion is the required heat energy when 1 mole of a substance as S.T.P. completely combusts in oxygen, forming it’s products at S.T.P.
• Methanol,
• Propanol,

                        -1409                                                +1371

Using Hess’ Law,

• Enthalpy of Solution

• This is the heat energy required when 1 mole of a substance dissolves in sufficient water that no further energy change occurs, at S.T.P.
• is the when 1 mole of a substance is made from its gaseous ions:

• is the when 1 mole of gaseous ions are completely hydrated:

Entropy Change,

• It is a measure of the number of possible ways that a system can be organised. This includes both the particles in the system and the quanta of energy within the system.
• For anything spontaneous, S must increase (positive)

• k → Boltzmann constant
• W N° of ways =
• Units →

Kinetics

Rate of Reaction

• The rate at which the concentration of a particular reactant decreases (as the concentration of the product increases, per unit time). Given in
• Ways of measurement:

• Change in mass as gas escapes.
• Collect a gas given off and look at volume changes.
• The formation of a p.p.t.
• Use of an indicator to show the end of reaction (different [A] and T)

Collision Theory

• For a reaction to occur:

• Particles have to collide with each other.
• Must collide with sufficient energy.

• → The minimum amount of energy required.
• Any factor that increases the frequencies of collisions or the energy with which they collide makes the reaction go faster.
• Reactions occur when reacting species have

The Rate Determining Step

• The slowest step in the chemical reaction (there are multiple steps).
• If we know the r.d.s. we can know what catalyst to use to speed up the reaction.
• If the reactant isn’t in the r.d.s then you get a linear relation between conc. and rate of rxn.

• Reaction order: the order of reaction with respect to a particular reagent.

• Temperature

• As T increases the particles will move faster to give more collisions per second.
• As T increases, the more of the particles will possess the necessary  (increased proportion of molecules with .

• Surface Area

• Increasing S.A. increases the area for collisions.
• In a solid substance only the particles on the surface can come into contact with a surrounding reactant.

• Concentration

• The more concentrated the reactants, the more collisions there will be per second per unit volume.
• As concentration decreases in reactants, the reaction slows down.
• The degree of change depends on whether a reagent is involved before, during or after the r.d.s.
• In concentration of reagent involved in r.d.s. is increases, the r.o.r. rises but it is not proportional.

• x→ reaction order with respect to A (depends on where A takes part in the reaction sequence, how many steps A takes part in and how many particles of A are in each step.
• [A] → concentration of A

Catalysts

• Increases the rate of reaction without being changed themselves.
• Works essentially by bringing the reactive parts of the reactant particles into contact with each other.
• Lowers .

Equilibrium

Dynamic Equilibrium

• Some reactions are reversible- they spontaneously go in both directions.
• Equilibrium can be approached from both directions.
• Equilibrium occurs where, rate of forward reaction = rate of backward reaction.
• Dynamic equilibrium → in a closed system the concentration of all the reactants and products will eventually become constant.

• Water (Phase Equilibrium)

• Rate of vaporization = Rate of condensation.
• Any liquid exists in equilibrium with its gas. These equilibria shift their position in exactly the same way as chemical equilibria.
• Heat shifts the equilibrium to the right; pressure, to the left.

The Equilibrium Constant

•  for a homogenous reaction.
•  or   
•  can remain the same when equilibrium shifts. This relies on the changes simply cancelling out.
• The magnitude of :

• Is related to the position of equilibrium.
• When >>1, the reaction goes almost to completion (products favoured)
• When <<1, the reaction hardly proceeds.

Le Chatelier’s Principle

• A system at equilibrium will shift the position of equilibrium to offset any changes that are made to the conditions of the system.
• E.g. If we raise the temperature then the system will favour a shift in the endothermic direction in order to reduce the temperature.

Factors Affecting the Position of Equilibrium

• Concentration

• If we increase the concentration of the reactants, the point of equilibrium shifts towards the products (right), to make more products. So, the concentration of the products increases while that of the reactants decreases.

• Temperature

• Exothermic Reactions → heat is also a product. So, taking heat away will shift the equilibrium to the right in order to make more products, therefore the forward reaction in exothermic processes is increased by lowering T.
• Endothermic Reactions → the exact opposite occurs, therefore these prefer higher temperatures, on order to make more products.
• Also affects . For exothermic reactions, a rise in T decreases the concentration of products, so the value of  decreases. The opposite is true for exothermic reactions.

• Pressure

• An overall change in gaseous reaction of volume occurs.
• Increasing the pressure will move the equilibrium towards the side with less volume. This shift reduces the total number of molecules in the equilibrium system and so tends to minimize pressure:

        → raise in P shifts the equilibrium to the right.

Catalysts

• Increases the rate at which equilibrium is reached. The point of equilibrium does not change.
• They affect backward and forward reactions equally.
• Changes the reaction pathway for the reactions in question.

The Häber Process (Production of Ammonia)

• needs catalyst → Fe or Pt
• High activation energy.
• Forward reaction is exothermic. Backward is endothermic.
• Forward reaction is initially very fact as there’s lots of starting material (high pressure or concentration). The backward reaction is initially very slow.
• At equilibrium, both reactions continue.
• Favours high pressures (fewer moles of gas on products side). Raise in P shifts equilibrium to the right.
• Favours low temperatures (forward reaction is exothermic). Raise in T shifts equilibrium to left, so more reactants are formed and gets smaller as the denominator increases.
• Industry→ pressure in the hundreds of atmospheres range (~250 atm) as the equipment provides high enough pressures without excessive costs or risk. Temperatures aren’t very high (350-400 oC) even if only 15-20% yield is obtained because lowering the temp slows the r.o.r. so it would take longer to reach equilibrium.

Catalyst process (Production of Sulphuric Acid)

• Sulphuric acid is used for fertilizers, paints, detergents, fibers, etc.
• Forward reaction is exothermic (= -197 kJ mol-1) so it is favoured by low temperature.
• Less moles of product is gas. So high pressure is favoured.
• Catalyst: Pt or Vanadium (V) oxide  
• T = 450 oC, P = 2 atm. Produces 99% yield, even with the low P.

Acids and Bases

Properties of Acids and Bases

• Acid: a substance that will give  ions in aqueous solution.
• Base: a substance that can neutralise an acid (accepts  ions)
• Alkali: a substance that will release ions into the solution. All alkalis are bases, but no all bases are alkalis.

• Acid neutralization with a base:

• Acid + Hydroxides→  salt + water
• Acid + Metal Oxide →  salt + water
• Acid + Ammonia →  salt  ()

• Acid neutralization with a metal:

• Acid + metal → salt + Hydrogen
• Acid + metal carbonate→ salt + water + carbon dioxide
• Acid + metal hydrogencarbonate→ salt + Hydrogen + carbon dioxide

()

• Acids:

1. Hydrogencarbonates are acidic salts:
2. Small charged cations (e.g. ) have acidic salts:
3. Hydrogen sulphates:
4. Ammonium salts:  

• Bases (unusual)

1. Ammonia:
2. Hydrogencarbonates:
3. Hydrogen sulphates:

From the above reactions we cans see that the product of an acid reaction acts as a base in the reverse reaction, and vice versa.

• Strong and Weak acids

• A strong acid is one which fully dissociates (ionizes) in aqueous solution. E.g. .
• A weak acid is one which only partially dissociates in aqueous solution. E.g. Ethanoic acid, carbonic acid:  
• In Water

• Strong acid        →
• Weak acid         →
• With the same concentrations, the pH of a weak acid is higher than that of a strong acid.

• Conductivity of a strong acid is higher because it has more ions.

• Strong and Weak bases

• A strong base is one which fully dissociates (ionizes) in aqueous solution. E.g. .
• A weak base is one which only partially dissociates in aqueous solution. E.g. , ethylamine, amines.
• In Water

• Strong base        →
• Weak base         →
• With the same concentrations, the pH of a strong base is higher than that of a weak base.

• Conductivity of a strong base is higher than that of a weak base.

The pH Scale

• Each change of one pH unit represents a tenfold change in the hydrogen ion concentration
• pH is equal to the power of 10 of the hydrogen ion concentration.
• pH is temperature dependant, it also has a scale of 1 to 14.
• Water:

• Only about of are present at 25 oC.
• For every we have a . When []=[] then the water is “neutral”
• Water will behave as both acid or base:

•         → acid
•                 → base

The Ionic Constant

Indicators

• All indicators are weak acids.
• Their dissociated form produces an anion that is different colour from the undissociated acid in aqueous form.
• Adding acids or alkali shifts the equilibrium and brings about colour change.
• At the equivalence point, roughly equal amounts of HI and I- are present.

Buffer Solutions

• Buffer solutions will maintain a specific pH despite the addition of small quantities of and .

• Ethanoic acid                        

Sodium Ethanoate                

• Adding Acid→ combine with large reservoir of from the NaA

•  will form mostly undissociated HA and are taken out.
• pH is held constant

• Adding alkali→will combine with from the weak acid.

• As are used up, more HA dissociates to restore equilibrium.
• are mopped up.

• Another example:                         
• Preparing Buffer Solutions:

• Acidic         →a weak acid and its salt, with a strong base.

→taking a solution of a strong base w/excess weak acid, to then leave the salt and unreactive weak acid:

• Alkaline        →a weak base and its salt, with a strong acid.

• Blood is an example of a complex buffer solution. It works within a narrow pH to allow the interaction of O2 and haemoglobin.

Acid-Base Titrations

                        14                                

                        12                                                       3

                         

                                                   2        

                                            7        

                         2        

                         1

        

1. Alkali begins to neutralise ions () We’re adding more so equilibrium moves to the right.
2. Rapid pH change. All the  from the acid is moped up at the bottom of the steep curve. Water produces few which are rapidly neutralised by adding  causing a rise in pH.
3. Very little remains. Adding is unlikely to produce much of a reaction so now pH increases slowly.

Oxidation and Reduction

Redox Reactions

• Oxidation→ is the loss of electrons in a substance.

• An oxidising agent is one that readily oxidises other substances. By doing so it gets reduced.

• Reduction→is the gain of electrons in a substance.

• An reducing agent is one that readily donates electrons (reducing them), thus oxidising itself.

• Half equations show either the reduction process or the oxidation process.
• How to put charges:

• Assume everything is purely ionic.
• Some things have only 1 non-zero oxidation state. Oxygen’s is 2-, except in peroxides, where it is 1-.

Oxidation Numbers and the Name of Compounds

• Oxidation numbers in names of compounds are represented in roman numerals.
• When elements show more than one oxidation state they have a number:

• Eg.         → Iron (II) chloride.

→ Potassium dichromate (VI)

Redox Titrations

and

1.                                 (oxidised)
2.         (reduced)

• N°  of moles =
• From V of titration calculate moles of  (known C)
• Use mole ratio to work out moles of  (unknown C)
• Work out C of using rearranged equation.
• Large quantities of required.
• Assume all oxygen turns to water.

Reactivity Series

• K, Na, Li, Ca, Mg, Al, C, Zn, Fe, Sn, Pb, H, Cu, Ag, Au, Pt
• The more readily a metal loses  the more reactive it is.
• Displacement reaction:

• The most reactive metal takes away from the least reactive.

A Voltaic Cell

• A Voltaic half-cell is a metal in contact with an aqueous solution of it’s own ions.
• A Voltaic cell consists of two different half-cells, connected together by an external wire and a salt bridge (damp filter paper soaked in). This allows electrons transferred during the redox reactions to produce electricity.
• E.M.F.→ electromotive force,

                                             

                           + 0.76 V                                        +0.34 V

• Zn is more reactive so it ionizes
• Cu electrode gets heavies as the gains electrons to become.
• For the cell :

• = 0.76 + 0.34 = 1.10 V (E.M.F. needs to be positive.

• Hydrogen Half-Cell

The p.d. obtained when an electrode of some substance is placed in a 1M solution of it’s ions to make a ½ cell and is used to complete a circuit using a hydrogen half-cell at S.T.P. → standard electrode potential.

Electrolysis

• It is the process of turning electrical energy into chemical energy.
• Works by redox reactions.
• A voltaic cell’s electricity is produced by the spontaneous redox reduction taking place. Electrolytic cells are used to make non-spontaneous redox reactions by providing electricity from an external source.
• And electrolyte is a substance which doesn’t conduct electricity when solid, but does when molten or dissolved in aqueous solution; and it’s chemically decomposed in the process.
• The flow of  makes up the circuit.
• Used to obtain the elements which are above Al in the reactivity series (the ones below are done in a blast furnace using C).

• The electrolysis of molten NaCl:

• Sodium is formed at the Cathode.
• Chlorine is formed at the Anode.

• The least reactive element is ALWAYS discharged.

Electroplating

• Eg. Copper plating.
• The metal to be plated must be the cathode.
• The other electrode (anode) is made from copper.
• As electricity flows, the Cu anode dissolves to form in solution. The of the are deposited onto the cathode.
• The metal is now covered in copper.

Organic Chemistry

Homologous Series

• A group of compounds with similar chemical properties, but different molecular formulae. E.g. Alkanes.
• Boiling Points → Inside a series, it increases as mass increases, since a higher mass means more electrons and therefore a possibility of higher Van der Waal’s forces.

Hydrocarbons

• Alkanes

• Simple  bonds; formula .
• Highly flammable.
• Relatively inert because their internal bonds are very strong, giving high activation energies. Also, the bonds are non-polar making them less vulnerable.
• Good solvents.
• Melt and boil easily (with steam) due to weak Van der Waal’s forces.
• Completely combust in sufficient oxygen, producing CO2 and H2O. This is exothermic.
• Example: Hexene

1. Molecular Formula        →        
2. Structural Formula        →        
3. Graphical Formula        →

• Alkenes

• Double bonds; formula .
• More reactive than alkenes because of the double bond.
• Test → Bromine water. Goes from brown to colourless.
• The  is much more prone to react, but the other simple bonds are just as inert as in alkanes.
• Can combust in sufficient oxygen, producing CO2 and H2O. May combust incompletely to make C (soot) and CO.

• Isomers

• Hydrocarbons of the same structural formula, but different structural formulae.

• E.g. Cyclohexane and Hexene →

• Isomers have equal numbers of electrons and therefore would be expected to have the same Van der Waal and b.p.t. However, they are more spherical in shape and this reduces surface area, therefore lowering both Van der Waal’s forces and b.p.t.

Stability of Carbon Chains vs Silicon Chains

• Si will also form chains with itself but they are not stable in an oxidising atmosphere.
• Stability comes from two places:

• C-C and C-H are similar in energy with C-O. So there is little to be gained in energy in oxidising a hydrocarbon. This is not true for Si.
• Si-Si and Si-H are much weaker than C-C and C-H. So the activation energy for combustion of silicon hydride organic analogues is so low that it happens at S.T.P.

Naming Convention (IUPAC)

1. Find the longest carbon chain (you can count around corners)
2. This will be named according to the sequence meth, eth, prop, but, etc.
3. Chose the principal functional group to get the ending, e.g. Alkanes end in –ane, alkenes in –ene, alcohols in –ol, etc.
4. Branching of the carbon chain:

• Decide on the length of the branch → 1C is methyl, 2Cs is ethyl, etc.
• Decide on the location of the branch:

• 2nd C atom on branch → 2-methyl.
• 3rd C  atom on branch → 3-methyl.

Make this number as small as possible.

5. Other substituents on the C chain in addition to branches:

• -Cl → chloro-
• -Br → bromo-
• -I  → iodo-
• -OH → hydroxyl- (except in alcohols)
• =O  → carbonyl- (except in acids, aldehides and ketones)

Reactions with Hydrocarbons

• Additions:

• Hydrogenation

• Ethanol         

• Bromoethanol         

• Polyethene

• Alcohols

   Ethanol (alcohol)                        Ethanal                        Ethanoic acid (carboxylic acid)

• Esters

                                                Ethyl ethanoate

• Structural Formulae Pattern                                
• R → Any alkyl group

R’→ Another alkyl group

• Forms Alkyl alkanoates