CHEM REVISION A2 NOTES PDF

CHEM REVISION A2 NOTES Chemical energetics The ΔHat is always endothermic as energy is always required to break any bonds between the atoms in the element to make them gaseous atoms -when ions are combined to form an ionic solid lattice there is an extremely large release of energy - This suggest...

CHEM REVISION A2 NOTES Chemical energetics The ΔHat is always endothermic as energy is always required to break any bonds between the atoms in the element to make them gaseous atoms -when ions are combined to form an ionic solid lattice there is an extremely large release of energy - This suggests that an ionic compound is much more stable than its gaseous ions due to the strong electrostatic forces of attraction between the oppositely charged ions in the solid lattice so the gaseous ions are less stable than the ions in the ionic lattice - The more exothermic the value is, the stronger the ionic bonds within the lattice are NB: when they just talk about lattice enthalpy without specifying which one it is referring to lattice formation All of these definitions are for standard conditions which are 101kPa and 298k (25 degrees) -Standard conditions and s.t.p. ARE NOT THE SAME s.d- 101kPa & 298k ,stp- 273k & 101kPa Factors affecting lattice enthalpy Ion charge -- increasing the ionic charge increases the attraction between the positive and negative ions meaning a larger, more negative lattice formation enthalpy. Ionic radius- decreasing the ionic radius means the ions are closer together in the lattice. Hence, the attraction between the ions is stronger meaning a larger, more negative lattice formation enthalpy. Both factors can be combined to make charge density so the greater the charge density the greater the enthalpy and vice versa Electron Affinity The first electron affinity (EA1) is the enthalpy change when 1 mole of electrons is added to 1 mole of gaseous atoms, to form 1 mole of gaseous ions each with a single negative charge under standard conditions Factors affecting electron affinity Nuclear charge: the greater the nuclear charge, the stronger the attractive forces between an incoming electron and the nucleus Distance: the greater the distance between the nucleus and the outermost shell/orbital where the electron is added, the weaker the force of attraction  Shielding: the greater the number of shells, the greater the shielding effect and the weaker the force of attraction Only the first EA is exothermic, so what does this mean? -The second and third electron affinities are endothermic, as energy is absorbed, this is because the incoming electron is added to an already negative ion so it is going to require energy to add the electrons mainly to electron repulsion hence, the energy is needed to overcome the repulsive forces between the incoming electron and negative ion Trends in electron affinity of Group 16 & Group 17 elements -Going down Group 16 and 17: The outermost electrons are held less tightly to the nucleus as they are further away The number of electron shells increases causing an increased shielding of the outermost electrons It gets more difficult to add an electron to the outer shell Less energy is released upon adding an electron to the outer shell So generally, the EA1 becomes less exothermic as you go down a group -Fluorine is an exception and has a lower EA1 than chlorine  Fluorine has a very small atomic radius, which means that the electron density of fluorine is high There is more repulsion between the incoming electron and the electrons that are already present in fluorine These repulsive forces reduce the attractive forces between the incoming electron and the nucleus As a result, the EA1 of fluorine is less exothermic than expected Enthalpy of solution and hydration Hydration enthalpy is basically a measure of the energy released when attractions are set up between positive or negative ions and water molecules. When an ionic solid dissolves in water, positive and negative ions are formed Water is a polar molecule with a δ- oxygen (O) atom and δ+ hydrogen (H) atoms which will form ion-dipole attractions with the ions present in the solution The size of the hydration enthalpy is governed by the amount of attraction between the ions and the water molecules: The attractions are stronger the smaller the ion. For example, hydration enthalpies fall (/become less exothermic) as you go down a group in the Periodic Table. The small lithium-ion has by far the highest hydration enthalpy in Group 1. The attractions are stronger the more highly charged the ion. For example, the hydration enthalpies of Group 2 ions (like Mg2+) are much higher than those of Group 1 ions (like Na+). Entropy - The entropy (S) of a given system is the number of possible arrangements of the particles and their energy in a given system,in other words, it is a measure of how disordered a system is  - When a system becomes more disordered, its entropy will increase, an increase in entropy means that the system becomes energetically more stable - The solid particles are more ordered in the solid lattice as they can only slightly vibrate When dissolved to form a dilute solution, the entropy increases as: The particles are more spread out There is an increase in the number of ways of arranging the energy The crystallisation of a salt from a solution is associated with a decrease in entropy The particles are spread out in solution but become more ordered in the solid Predicting entropy Gibbs free energy Since the feasibility of a reaction does not only depend on the entropy change of the reaction but can also be affected by the enthalpy change The Gibbs free energy (G) is the energy change that takes into account both the entropy change of a reaction and the enthalpy change LOOK at the units ……ALWAYS LOOK AT THE UNITS AND CONVERT SO THAT EVERYTHING IS IN THE SAME UNIT !!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!  Using Gibbs free energy equation to find other things You use these when you are trying to find what temperature a reaction for neg-neg & pos-pos is feasible, so you assume gibbs free=0 Equilibria - A conjugate acid-base pair refers to a pair of compounds that differ by one proton  pH - Stands for power of hydrogen For a weak acid Can ONLY use for weak acids because they dont disociate fully Acid dissociation constant(Ka) The acid dissociation constant (Ka) is a measure of the extent to which an acid dissociates in solution and therefore its strength. The less an acid dissociates, the smaller the value of Ka. The stronger the acid, the higher the value of Ka. The value of Ka is constant at a specific temperature. The equation for the dissociation of ethanoic acid in water -The equilibrium for weak acids lies far on the left so there are little ions. At any one time, only about 1% of the ethanoic acid molecules have converted into ions. The rest remain as simple ethanoic acid molecules. -Most organic acids are weak. Hydrogen fluoride (dissolving in water to produce hydrofluoric acid) is a weak inorganic acid -If HA is any given acid, the ionic equation for forming an acid is: General equation for the weak acid dissociation in equilibrium pKa Buffers - A buffer is a solution that resists change in pH when small quantities of an acid or alkali are added to it - →How is the buffer action of a solution related to the dissociation constant (Ka) of the acid used in the buffer? A buffer with a high Ka value will have a strong buffer action, meaning that it can resist changes in pH better than a buffer with a low Ka value. Acidic Buffers - An acidic buffer is made by adding a weak acid to a solution of one of its salts (its conjugate base) like for example, ethanoic acid and sodium ethanoate - In the first equation when H+ increases(basically adding an acid), the equilibrium will shift to the left to decrease H+ by moving to the reactant side so it can make more undissociated acid molecules. It can do this because there are lots of ch3coo- ions to react to form the acid. - When a base is added, the alkali accepts the H + ions, therefore decreasing the amount of H+ ions, so the equilibrium will shift to the right to produce more H+ ions. It  can do this because there are lots of undissociated acid molecules from the weak acid. Basic Buffers They are made by adding a weak base to their salt. For example, ammonia solution and ammonium chloride act as a basic buffer. - When H⁺ ions are added to the basic buffer(when an acid is added), they combine with OH ions released when the ammonia molecules partially dissociate. The equilibrium shifts to the right to replace the OH⁻ ions. - When a base is added the equilibrium will shift to the left to remove the OH- ions and turn them into water. They combine with NH₄⁺ ions from the salt, to make more undissociated ammonia molecules. Calculating pH of Acidic Buffer Solutions Henderson-Hasselbach equation Only use for buffers and weak acids Amphoteric buffers The amino group is alkali so it acts as a base The carboxyl group is acidic so it acts as an acid So when it reacts with an acid- H2NCH(R)COOH+ H+ —> H3NCH(R)COOH+ (the proton joins the amine group) So when it reacts with a base- H2NCH(R)COOH + OH- —> H2NCH(R)COO- +H2O (the hydroxide ion takes the H from the carboxyl group to make water) Buffer Action in Blood Human blood is kept at a pH between 7.35 and 7.45 by the buffer action of carbonic acid (H₂CO₃) and bicarbonate anions (HCO₃⁻ ). Two equilibriums are set up: Ionic product of water(kw) NB:You need to know the simplified version not be writing the ultra simplified version Since the value of of the ionic product of water is 1*10^-14 , That is the the highest pH value is 14 Solubility product - Solubility is defined as the number of grams or moles of compound needed to saturate 100 g of water, or it can also be defined in terms of 1 kg of water, at a given temperature  - Ksp is ONLY useful for sparingly solublesalts due to making an assumption that the concentration of C for the example (the salt) stays the same since the equilibrium lies far left - As Ksp increases the solubility ofthe salt also increases Calculating the solubility product from solubility values 1. Ensure the values are in moldm^-3 2. Write the equation 3. Write the Ksp equation 4. Work out the concentration of the product based of coefficients 5. Substitute values into expression Common ion effect - A saturated solution is a solution that contains the maximum amount of dissolved salt at that specific temperature  - If a second compound which has an ion in common with the dissolved salt, is added to the saturated solution, the solubility of the salt reduces and a solid precipitate will be formed. This is also known as the common ion effect. - The solubility product can be used to predict whether a precipitate will form or not. A precipitate will form if the product of the ion concentrations is greater than the solubility product (Ksp). If the product of the ion concentration is less than the Ksp then a precipitate will not form How to prove whether a precipitate will form or not: 1. Determine the equilibrium reaction for the saturated solution Eg CaSO4⇋CA2+ + SO42- 2. Write down the equilibrium expression for Ksp 3. Determine the concentration of the ions (keep in mind the new volume) 4. Substitute the value into the expression 5. Compare the value gotten with the Ksp. If the product of ion concentration is less than Ksp, a precipitate is not formed. If it is greater than Ksp, a precipitate is formed. Partition Coefficients The partition coefficient (Kpc) is the ratio of the concentrations of a solute in two different immiscible solvents in contact with each other when equilibrium has been established (at a particular temperature) Calculating the partition coefficient  Factors Affecting the Partition Coefficient Electrochemistry