Chemistry Preliminary Notes PDF

Module 1 - Properties and structure of matter Elements are pure substances made up of one type of atom ​that can’t be broken down into simpler substances. Compounds are pure substances made up of 2 or more different types of atoms​ and can be decomposed by heat, electrolysis, and high energy EM radi...

Module 1 - Properties and structure of matter Elements are pure substances made up of one type of atom ​that can’t be broken down into simpler substances. Compounds are pure substances made up of 2 or more different types of atoms​ and can be decomposed by heat, electrolysis, and high energy EM radiation. Compounds can exist in 3 ways: Ionic ​compounds (salts) - A metallic element bonded to one or more non-metallic elements forming a crystal lattice through a complete transfer of electrons. They don’t exist as discrete molecules but in a ratio of atoms. Covalent​ compounds - A group of 2 or more non-metal atoms bonded by sharing electrons to form discrete molecules. Lattice/Network ​- A 3D network of atoms chemically bonded in a lattice that are not discrete molecules. Mixtures are composed of non-chemically bonded substances​ and can either be: Homogeneous - Components are uniformly mixed​ and have one phase e.g. solutions, alloys, unsaturated salt water. ○ Alloys: Mixture of two or more elements with at least one of them being a metal. Alloys have different properties to the metals which they contain. This makes them more useful than pure metals. They are often harder than the metals they contain e.g. steel, brass, bronze, stainless steel. Heterogeneous - Not uniformly mixed​ and can have more than one phase e.g. muddy water, sand, rocks. True solutions: made where a solute is dissolved in a solvent.​ Below 10−7 m particle size Suspensions: where the solute does not dissolve in the solvent.​ The solute particles will settle and can be filtered out e.g. C aCO3 in water. Above 10−5 m particle size Colloids: particles are dispersed throughout but are not heavy enough to settle​. The particles are not as small as a solution but are smaller than those of a suspension. They can be separated by filtering e.g. milk, mayonnaise, paint, fog. 10−5 to 10−7 m particle size Physical changes only affect the form ​of the substance Substance changes its physical appearance but not composition Changes are reversible Involves the breaking of intermolecular forces but not bonds Particles stay the same but move differently A small amount of heat or energy may be taken in or given out Chemical changes affect the chemical composition​ of the substance A new substance is formed A gas may be given off There may be a change in colour There may be a change in temperature A precipitate may form Changes are irreversible STP: 0C 100kPa and SATP / RTP / NTP / SLC: 25C 100kPa unless given otherwise Mixtures in the Spheres of the Earth There are 4 spheres: Biosphere - Living organisms ○ Mixtures of carbon containing compounds Lithosphere - Earth’s outer mantle and crust ○ Contains mixtures such as soils, rocks, sand, mineral ores, coal, oil and natural gas. Hydrosphere - Earth's Water ○ Mixtures of water, sodium chloride, magnesium and calcium chlorides Atmosphere - Gases above Earth’s surface ○ Mixtures of gases e.g. N 2 , O2 , CO2 , H 2 O Separation techniques and their properties: Separation Method Property which enables separation Sieving Particle size Evaporation Liquid has a lower boiling point than solids Distillation Relatively large differences in boiling point between two substances Fractional Distillation Small differences in boiling point between two substances usually in the same phase Separating funnel Components are immiscible and have different densities Filtration One solid, the other a liquid or solution Sedimentation Density Qualitative​ analysis - Determining ​what​ substances are present Quantitative​ analysis - Determine ​how much​ of each substance is present To calculate the percentage composition by mass or calculate the mass of each component use the formulae: M ass of component T otal mass of mixture × 100 % of component × mass of the mixture Orbitals are geometric areas of 95% electron presence probability​ as determined by Schrodinger’s equation in which one or two electrons can move. Principal energy levels (1, 2, 3…) are groupings of similar energy sublevels (s, p, d, f) which are in turn groupings of individual orbital lobes. In order to have probability in excess of 99%, an area the size of the observable universe is required  The Aufbau principle governs that electrons fill​ atomic orbitals of the lowest available energy levels before occupying higher levels Exceptions to Aufbau principle occur in the d block due to the reduced energy distinction between s and d levels and Coulombic repulsion (incl. Cu, Ag, Cr) The Heisenburg uncertainty principle states that the position and momentum of a particle (such as an electron) can not be measured accurately simultaneously The Pauli exclusion principle and Hund’s rule states that ​orbitals of an energy level are always filled first by electrons with + 21 spin and are then completed with electrons of - 21 spin​ (identical fermions cannot occupy the same quantum state within a system). E.g. Oxygen Valence electrons are the electrons present in the outermost shells of an atom Elements in the same group have the same number of valence electrons Valency is the maximum number of direct bonds formable with the element. Variable valency may occur when an atom loses or gains more electrons than the valence shell. Orbital notation can be determined by filling to the appropriate orbital ​given the number of electrons or through the use of spdf blocks on the periodic table by counting the number of filled orbitals. Helium and groups 1 and 2 comprise s orbitals from energy levels 1-7 Transition metals comprise d orbitals from energy levels 3-6  Excluding helium, groups 13 - 18 comprise p orbitals from energy levels 2-7 Lanthanides and actinides comprise group f from energy levels 4-5 Atomic emission spectra and flame tests are based on the excitation of electrons: When electrons are supplied with energy they can become excited. De-excitation causes the release of energy. The re-emission of energy occurs as different wavelengths of light. Flame test: visible spectrum of light indicates ion Mass spectrometry analyses the spectrum of light released by means of dispersion through an electromagnetic lens and detection by an amplifier wherein positive species are detected and neutral species are eliminated Isotopes and radioactive chemistry: Isotopes are elements with the same number of protons and different numbers of neutrons and are named by their mass number. Given the notation ZA M. A represents the mass number and Z represents the atomic number. Ionising radiation is classified into several primary types: 4 Alpha radiation consists of the emission of a particle of 2 He with charge +2 and a large relative mass and reduced penetration (typically blocked by paper) Beta radiation consists of the emission of electrons ( β − , −1 0 e ) or positrons ( β + , 0 +1 e )​, charge is -1, mass is minimal, and penetration is moderate (blocked by sheets of lead or aluminium) Gamma and X-ray radiation consists of high energy electromagnetic radiation​, with charge 0 and mass 0, penetration is high (blocked by thick lead or concrete) Neutron radiation is able to penetrate through layers of concrete but decays quickly, has relatively increased mass, no charge, and high penetration (shielding typically consists of several metres of concrete or boron/cadmium compounds (dissolved in water) or metal) To calculate relative atomic mass (RMS) use: RM S = ∑(relative abundance of isotope × mass of isotope) Technetium 99m in medicine: 98 It is formed by 42 M o + 01 n → 99 42 M o → 99 43 T c + 0 −1 e Technetium 99m is used as a radiotracer in medical diagnostic procedures such as SPECT scans in oncological, cardiothoracic, and other fields Band of Stability: Periodicity: Increasing from the given corner: Upper-right → lower left: ○ Atomic radius ○ Metallic character Lower-left → upper right: ○ Ionisation energy​ - The energy needed to remove an electron from a ​gaseous ​atom ○ Electron affinity ○ Electronegativity​ - The ability of an atom to attract electrons ○ Non-metallic character Additionally:. Core Charge (The number of valence electrons) increases across a period  Reactivity increases from the upper right corner of s, d, f-blocks and from the lower left corner in p block, excepting noble gases Shielding increases​ with period Melting point and boiling point ​increase until across the period from group 1 to group 14, then drastically drops for groups 15-18. M.p. and B.p. decrease down a group, suggesting that metallic bonds in crystals of metals are weaker as atoms get larger. Valency +1 for group 1, +2 for group 2, variable for d-block, and +3, -4, -3, -2, -1, 0 for p-block ○ 1,2,3,t,4,3,2,1,0 Chemical bonding and intermolecular forces: Allotropes are different physical forms of an element in which an element can exist in a given phase. Example: Carbon has multiple allotropes including: diamond, graphite and fullerenes: Diamond: ○ Has a high refractive index ○ Hardest natural material - used in industrial cutting, drilling equipment ○ 3D strong covalent network lattice - strong, hard, rigid ○ Every carbon is bonded to 4 other carbon atoms in tetrahedral arrangement ○ Does Not conduct electricity ○ Brittle due to rigid structure ○ As atoms are held together very strongly, it is a good thermal conductor Graphite ○ Carbon atoms are bonded in flat hexagonal rings in layers, but weak dispersion forces between the layers make graphite hard in one direction but slippery and soft in another direction. ○ It is a dry lubricant ○ It is a good heat insulator ○ Each carbon is bonded to 3 other carbons, leaving a free valence electron. This delocalised electron is free to move, so graphite is a good conductor of electricity Fullerenes ○ Made up of molecules containing carbon atoms arranged in a cage like structure ○ Similarly to graphite, it can conduct electricity ○ They can be spherical or cylindrical ○ Can be used in drug delivery systems in the body Polarity and Intermolecular Forces: Polarity means there is a net difference in charge between two areas in a molecule​: If two bonded atoms are different, the electrons will stay closer to the more electronegative atom (has stronger pull). This produces a small charge imbalance, making one end of the bond more negatively charged δ − than the other δ +. The difference in electronegativity can lead to three outcomes: No difference → Nonpolar covalent Intermediate difference → Polar covalent  Very different → Ionic Polarity is determined by electronegativity and molecular geometry (VSEPR) Types of intermolecular forces list by strength in descending order: Ionic bonding​ commonly occurs between ​negatively and positively charged ions​ (sodium chloride crystals) Ion-dipole interactions​ occur between ​polar molecules and charged ions​ (sodium ions in water) Hydrogen bonding occurs when electropositivity induced (from F, N, O) on a hydrogen ion leads to dipole-dipole-type interactions​.For hydrogen bonding to occur, a hydrogen atom must be bonded to three very electronegative atoms which are F, N and O. This creates a large partial positive charge on the hydrogen and is attracted to the lone pairs on nitrogen, oxygen or fluorine. Hydrogen bonds are 10 times stronger than dipole-dipole bonds but only 1/10th the strength of an ionic or covalent bond. Hydrogen bonding is present in compounds such as water and ethanoic acid. The density of ice is less than liquid water as when water freezes, it forms a lattice, with molecules being held further apart than in liquid. Dipole-dipole interactions occur in molecules that have a net dipole​ such as in polar molecules (NH3 - NH3, ICl - ICl). These forces result from the electrostatic attraction between the positive and negative ends of two or more dipoles. These forces are relatively weak as the partial charges are small. However, the more polar a molecule is, the stronger the dipole-dipole forces will be. Ion-induced dipole interactions occur when a relatively stronger ion induces the polarisation of an electron cloud in another molecule (Fe2+ ion - dioxygen) Dipole-induced dipole interactions occur where a polar molecule with dipole moment asymmetricality induces the polarisation of an electron cloud in another molecule (HCl - dichloride) Dispersion forces occur when momentary dipoles are formed. Weakest of all IMFs. This temporary charge can influence a change in the charge distribution in nearby molecules. This creates a temporary charge between the negative end and a positive end of two different molecules. There are billions of these interactions being made and broken every second. The strength of a dispersion force increases as the size of the molecule increases as more electrons will increase the probability of temporary dipoles. Types of molecular structures: Crystalline structures are usually solid, non-volatile, and have high melting points: An attraction and combination between a cation and an anion will form an ​ionic lattice network​. They are large structures that naturally exist due to intermolecular electrostatic attractions. They conduct electricity when molten or aqueous as for substances to conduct electricity, they must have charged particles that are free to move. Ionic compounds are hard. There are no discrete molecules and their formulas are empirical.  Metallic networks​ ​have a sea of delocalised electrons surrounding a positively charged cation with electrostatic bonding and good conductivity of electricity and heat. The electrostatic attraction between the cations and delocalised electrons is called ​metallic bonding​. They have lustre and are malleable. E.g. zinc and sodium Covalent lattice networks​ have covalent bonding throughout the structure and are similar to ionic lattice but have higher melting points and generally do not conduct electricity e.g. Diamond, Graphite, Silicon Dioxide, Silicon Carbide. They are of tetrahedral shape. Covalent molecular compounds ​are usually amorphous liquids or gases, volatile, have low boiling points, are poor conductors, and are soft. Molecular geometry: The octet rule states that ​atoms generally combine to obtain the noble gas configuration (8 valence electrons)​. A share of 8 valence electrons is the most stable arrangement of electrons, however it doesn’t apply to certain molecules including B F 3 , P Cl5 and SF 6. VSEPR is a theory used to predict the geometries of molecules with a central atom around which bonding occurs by modelling the repulsion of valence shell electrons from each other Tetrahedral​: ○ This is the most stable way for four equal electron clouds to arrange themselves ○ There are 4 bonding pairs, and no lone pairs ​[4,0] ○ 109.5 degree angle between atoms ○ E.g. N H 4 , CH 4 , SiCl4 , CCl4  Trigonal pyramidal​: ○ Tetrahedral arrangement of electron pairs without one of the atoms. ○ Has 3 bonding pairs and 1 lone pair ​[3,1] ○ 107 degree angle between atoms ○ E.g. N H 3 , P Cl3 Bent shape​: ○ Tetrahedral arrangement of electron pairs but only two are bonded. ○ 2 bonding pairs and 2 lone pairs ​[2,2] ○ 105 degree angle between atoms ○ E.g. H 2 O, H 2 S, CH 2 Linear: ○ Tetrahedral arrangement of electron pairs but only one bonding pair, or 2 bonding pairs with no lone pairs ( C O2 , HCN , BeH 2 ) ​[2,0] [2,3] ○ 180 degree angle ○ E.g. H Cl, HF  Trigonal planar: ○ Repulsion between electron pairs is minimised by a triangular arrangement about the central atom such that the electron pairs are in the same place. ○ 3 bonding pairs and no lone pairs ​[3,0] ○ 120 degree angles between atoms ○ E.g. B H 3 , BF 3 , SO3 , N O−3 A - Central atom X - Ligands to other atoms E - Lone pairs Linear: [2,0] [2,3] (CO2) Bent: [2,1] [2,2] (H2O )Trigonal pyramidal: [3,1] T-shaped: [3,2] Tetrahedral: [4,0] Disphenoidal:[4,1] Square planar:[4,2]Trigonal bipyramidal:[5,0] Octahedral: [6,0] (SF6) Lewis diagrams: Dots on lewis diagrams represent the number of valence electrons in the given atom (octet rule generally states a full shell is 8 valence electrons) For ionic compounds, bracket the ion structure with the charge Line bonds represent covalent bonds only  Democritus: Shit be made of atoms Dalton: Different elements have atoms of differing size and mass Thomson: Plum pudding model: small electrons in a large area of positive charge Rutherford: Atom has a positively charged nucleus (gold foil experiment); electrons orbit the nucleus Bohr: Electrons exist in discrete energy levels Schrodinger: Quantum/geometric model of electron position (orbitals) Chadwick: Discovery of neutrons in the nucleus Module 2 - Quantitative chemistry Stoichiometry Mass, energy, and number of atoms are always conserved in reactions. Balanced chemical equations show mole ratios. Relative atomic mass of an element is calculated by the weighted average of isotopic masses (protons and neutrons) m To calculate ​density​ , use the formula​ d= v. As temperature increases, density decreases. Relative molecular mass (molecular weight) is given by the sum of the relative atomic masses of the atoms as given in the molecular formula. M ass(g) m N umber of particles The ​number of moles​ (n) can be found with​ n= M olar M ass = MM = NA Number of atoms in a mole is given by the Avogadro constant (number of atoms in 12g of C-12): N A = 6.022 × 1023. To find the number of atoms, use N = n × N A. The volume of 1 mol of an ideal gas​ at 25C and 100kPa (SATP) is 24.79 L and is 22.71L at 0C and 100kPa. Molecular or formula mass is given as g mol−1 To calculate the ​empirical formula​ from experimental data in a given compound: find the moles of each compound and find the relevant ratio of moles. In a reaction, the ​Limiting reagent​ is a substance which does not contain enough moles for the complete reaction of another substance to occur, resulting in a leftover reactant. Preparing solutions Concentration is measured in three ways: Percentage by volume (%v/v), where both volumes are the same unit e.g. (mL/100mL). Percentage by mass (%w/w) where both masses are in the same unit e.g. (g/100g). Percentage by mass/volume (%w/v) where units are preferably (g/100mL) or (kg/100L). Some conversions: ppm = mg/L = mg/kg = mL/1000L = g/1000kg 1g = 1000mg 1mg = 1000𝜇g (micrograms) 1ppm = 1𝜇g/g = 1mg/kg 1ppb = 1𝜇g/kg = 1mg/tonne n Concentration​ of a solution in mol/L is given by:​ c = v or c = M Mm×v where v is in litres and mass is in grams. Hence, the number of moles can also be calculated by n = c × v Dilution​ of a given solution by the same solvent is given by​ c1 v 1 = c2 v 2 −1 Additionally: 1 PPM = 1 mg L = 1 g m−3 = 1 g / 1000L Measuring cylinders are used for approximate measurements of about ±5%, pipettes and burettes deliver variable accurate volumes of about ±0.2% to ±0.5%, whilst volumetric flask deliver fixed volumes with the same accuracy as the latter. Preparing a standard solution: 1. Calculate the mass of a substance you need to dissolve in a certain volume to make a desired m concentration for a solution using c= M M ×v 2. Weight the desired mass and transfer to a clean beaker. 3. Add distilled water until the mass is fully dissolved. 4. Transfer the solution to a desired volume volumetric flask using a filter funnel. 5. Rinse the beaker while transferring the washings into the volumetric flask. 6. Add distilled water until the bottom of the meniscus is on the calibration line. 7. Cork the volumetric flask and shake to ensure it is completely dissolved. Gas laws: Kinetic molecular theory of gases states that gases are made up of atoms or molecules, and as the temperature rises, so does the kinetic energy of the gas particles. F orce P ressure = 2 Area where force is in Newtons and area is in m. Units of gas pressures:  1n/m2 = 1P a (pascal). 1000Pa = 1kPa 1atm (atmosphere) = 101.3kPa 1bar = 100kPa Boyle’s law:​ ​Pressure is inversely proportional to Volume​. Given by the formula: P 1V 1 = P 2V 2 V1 V2 Charles’s law​: ​Volume is proportional to Temperature​ given by the formula:​ T = T 2 where 1 temperature is in Kelvin. PV Gay-Lussac’s gas law:​ ​Pressure is proportional to temperature​ Combined gas law: T is constant​. Also P1 P2 given by the formula​ T = T where temperature is in Kelvin. 1 2 Gay Lussac’s law of combining volumes:​ When measured at ​constant temp and pressure, volumes of gases in a reaction show simple whole number ratios​ to each other given by the formula: P 1V 1 P 2V 2 T 1 = T 2 where temperature is in Kelvin. Avogadro’s Law: States that equal volumes of has at the same pressure and temperature contain an equal V V2 number of molecules given by the formulae: V = N A × n and​ n 1 = n2 1 Ideal gas law:​ P V = nRT where r = 8.314. Non-ideal gas behaviour occurs at high pressures (ideal gas laws do not account for space taken up by the molecules themselves) and low temperatures (intermolecular attractions result in decreased pressure) For stoichiometric calculations use a BCA table. Module 3 - Reactive chemistry Types of reactions: Synthesis: reactions in which two or more substances combine to form a new substance: ○ A + B → AB Decomposition: When a substance decomposes into two or more substances by light, heat or electricity: ○ AB → A + B Acid-base (neutralisation) reactions: acid reacts with a base to form a salt and water depending on the reaction. ○ E.g. H Cl(aq) + N aOH (s) → N aCl(aq) + H 2 O(l) Combustion reactions: When elements or compounds are burnt in oxygen. Complete combustion will form carbon dioxide and water while incomplete combustion can form carbon or carbon monoxide: ○ Complete Combustion e.g: C H 4(g) + 2O2(g) → C O2(g) + 2H 2 O (l) ○ Incomplete combustion e.g: C H 4(g) + O2(g) → C (s) + 2H 2 O (l) , C H 4(g) + 23 O2(g) → C O (g) + 2H 2 O (l)  Precipitation reactions (double displacement): When two solutions of soluble ionic compounds are mixed together, they can either mix together completely or form an insoluble precipitate (ppt). ○ AB + C D → AD + C B ○ Precipitation e.g: AgN O3(aq) + N aCl (aq) → N aN O3(aq) + AgCl (s) ○ Solubility laws for precipitates: Acid-carbonate reactions: Acid reacts with carbonate to form water, salt, and carbon dioxide ○ E.g. N a2 CO3(s) + 2HCl(aq) → 2N aCl(aq) + H 2 O(l) + C O2(g) Metal - Oxygen reactions: A metal burns in oxygen to form a metal oxide: ○ E.g. 2M g (s) + O2(g) → 2M gO(s) Metal - Water reactions: Metal reacts with water to form a metal hydroxide and hydrogen gas: ○ E.g. 2N a(s) + 2H 2 O(l) → 2N aOH (s) + H 2(g) Acid-metal reactions: acid reacts with a metal to form a metal salt and hydrogen gas ○ E.g. Z n(s) + H 2 SO4(aq) → Z nSO4(aq) + H 2(g) Metal Displacement Reactions: The more reactive metal will displace a less reactive metal in a compound: ○ A + BC → AC + B ○ E.g. Z n(s) + C uSO4(aq) → Z nSO4(aq) + C u(s) ○ Determine which metal will displace another through the metal activity series or from highest to lowest on the reduction potentials table.  ○ Metals with a lower ionisation energy, large atomic radius and lower electronegativities will be the most reactive. ○ The metal activity series can be seen below: Reduction-oxidation (redox) reactions: Oxidation refers to loss of electrons and reduction refers to the gain of electrons. A reaction is a redox reaction if there is a change in oxidation number, which can be found using the following table: Rule Examples (O.N is in exponent) 1. The oxidation number (O.N) of a free N a0 , C 0 , Cl02 , P 04 element is 0. 2. The oxidation number of a simple ion is N a+1 , Cl−1 , M g +2 , Al+3 , N −3 equal to the charge of the ion. 3. Main group elements have an oxidation Ionic compounds: number equal to their charge on ions. K +1 Cl, M g +2 SO4  4. Hydrogen in a compound has an oxidation Compounds of H: H 2+1 O number of +1 except for metal hydrides where Metal hydrides: N aH −1 , CaH 2−1 it has an oxidation number of -1. 5. Oxygen has an oxidation number of -2 Compounds of O: H 2 O−2 except for: Peroxides: H 2 O2−1 , BaO2−1 In fluorine it has an O.N of +2 In peroxides it has an O.N of -1 6. The sum of the oxidation numbers in a C +4 O2−2 neutral compound is 0. Note: + 4 + 2 ×− 2 = 0 7. The sum of oxidation numbers in a In the polyatomic ion S O42− , S has polyatomic ion is equal to the charge on the O.N of +6 and O has an O.N of -2. ion. + 6 + 4 ×− 2 =− 2 8. The most electronegative element is given O+2 F 2−1 the negative oxidation number. Redox reactions can be shown in: ○ Half equations showing the transfer of electrons ○ Neutral species equation: shows the compounds involved in the reaction ○ Complete ionic equation: shows all ions present in the reaction mixture ○ Net ionic equation: shows the ions undergoing the reaction omitting spectator ions Deriving the different types of equations: ○ Balanced formula (neutral species) equation: 2KI (aq) + P b(N O2 )3 (aq) → P BI 2 (s) + 2KN O3 (aq) ○ Total ionic equation (also complete ionic): 2K + (aq) + 2I − (aq) + P b 2+ (aq) + 2N O3 − (aq) → P BI 2 (s) + 2K + (aq) + 2N O3 − (aq) ○ Net ionic equation P b2+ (aq) + 2I − (aq) → P bI 2 (s) Galvanic cells and standard electrode potentials Galvanic cells generate electricity through reduction-oxidation reactions Anode: Oxidation occurs to produce free electrons in the more reactive metal. Cathode: ​Electrons flow to the cathode​ and reduction occurs to create positively charged ions Salt bridge: completes the circuit and conserves solution neutrality. Without the salt bridge, a build-up of charge in the solutions and this would prevent further electrons or ions from forming due to chemical equilibrium. The ​anions flow towards the anode​ to balance the positive ions and the ​cations flow towards the cathode​ to balance the negative charges. The standard hydrogen half-cell consists of platinum metal immersed in a 1 mol/L solution of H+, with hydrogen gas being bubble over the electrode at 100kPa For reactions occurring in the direction of reduction ( K + + e− → K (s) ): a higher value indicates a stronger tendency to form the non-ionic form whilst a lower value indicates a tendency to form ions Voltage of a galvanic cell under standard conditions is given by: o o o E total = E reduction + E oxidation Galvanic cells are ​spontaneous when the cell voltage is positive Oxidising agents themselves are reduced and vice versa Galvanic cell notation X (s) | X + (aq) || Y + (aq) | Y (s) where the left is the anode / oxidation Example of a Galvanic Cell: Copper is gaining electrons, and hence the cations attach themselves to the cathode and become neutral atoms -> the cathode becomes bigger Zinc is losing electrons and the neutral atoms are becoming ions dissolved in the water -> the anode becomes smaller Electron loss in a reaction is denoted on the RHS, whereas gain is on the LHS Z n (s) | Zn2+ (aq) (1M ) || Cu2+ (aq) (1M ) | Cu (s) for the given example Rates of reactions Collision theory:​ States that reactions occur as a result of collisions between the reacting particles. It is based on the kinetic energy of particles. For a reaction to occur, the reactant particles must: Collide with each other Collide with sufficient energy to break the bonds within the reactants and reach the activation energy Collide with the correct orientation to break the bonds within reactants to allow for the formation of new products Activation energy:​ The minimum energy that a collision must possess for a reaction to occur. The rate of reaction can be measured through ​time ​or ​loss of mass​. Factors that change the rate of reaction include: Surface area - Higher surface area provides more surface for collisions to occur. Temperature - Higher temp increases kinetic energy, increasing the probability of collisions. Concentration - Higher concentration means there are more particles, increasing the probability of collisions. Pressure (gases) - Higher pressure moves particles closer together, increasing the probability of collisions. Presence of a catalyst. Catalysts speed up the rate of reaction but are not used up themselves. They provide an alternative pathway for a reaction which has a lower activation energy. Due to  this, more particles will be able to collide with sufficient energy for a reaction to occur Examples are: ○ Iron (Fe) - used in the production of ammonia N 2(g) + 3H 2(g) → 2N H 3(g) ΔH =− 92kJ/mol ○ Platinum (Pt) - Used in cars to catalyse carbon monoxide into carbon dioxide. C O(g) + 21 O2(g) → C O2(g) ○ Potassium permanganate ( K M nO4 ) - Used to decompose hydrogen peroxide into water and oxygen. 2H 2 O2(aq) → 2H 2 O (l) + O2(g) ○ Biological catalysts (enzymes) - Chlorophyll in leaves increases the rate of reaction of photosynthesis. Heterogeneous catalysts are of a different phase to reactants (example: gas - liquid) Homogeneous catalysts are of the same phase to reactants. Homogeneous reactions occur in a single phase (example: liquid - liquid) Module 4 - Drivers of reactions All chemical substances hold chemical potential energy. Bond Breaking requires energy​ while ​bond making releases energy​. Exothermic reactions release heat into the surroundings whilst endothermic reactions absorb heat into the system: In ​exothermic​ reactions, the energy released when making bonds is greater than the energy needed to break bonds. Endothermic​ reactions require a constant input of heat and cannot be self-sustaining i.e. the energy required to break bonds is larger than the energy released when making bonds. Enthalpy​ represents the total heat content of a system and is given in joules (J)  Change in enthalpy is given as ΔH. It can be calculated using the formula: ΔH = H products − H reactants where H is the heat content of a substance. Heat of combustion​ ( ΔH c ) is the energy released when 1 mol of a substance combusts completely at SLC. To determine energy released when n mole of a fuel burns use: E nergy released = n × ΔH c when ΔH c is given. The determine the heat of combustion from an experiment, use: q = mcΔT where q is the amount of energy in joules, c is the specific heat capacity (usually water (4.18)), and ΔT is the change in temperature. Specific heat capacity​ is the amount of energy (joules) needed to increase 1g of a substance by 1K. The specific heat capacity of water is very high (4.18), meaning it takes more energy to raise the temperature of water than other substances. The molar heat of combustion​ ( ΔH ) is the heat energy absorbed/released per mole of that substance. It is q calculated by:​ ΔH = where q is in kJ and ΔH is in kJ/mol. If a reaction is exothermic, ΔH < 0 n and if endothermic, ΔH > 0. Calorimetry​ is the process by which heat transfer is measured. The efficiency of calorimetry is a measure of how efficiently heat energy is transferred from the combustion of a material to the heating of water. It expected energy absorbed can be calculated by:​ % ef f iciency = actual energy supplied × 100. The expected energy absorbed will always be lower than actual energy supplied as heat is always lost to surroundings. Dissolution of an ionic lattice in water: When compounds are added to water, the positive ends of the water molecule are attracted to the anion and the negatively charged end is attracted to the cation. This is an ​ion-dipole attraction​. If the ion-dipole attractions are strong enough, the ions can be separated and become surrounded by water molecules. These are said to be ​hydrated​. The process of separating ions from a lattice is called ​dissociation​. There are ​three steps in dissolving​: Ionic bonds with the crystal lattice are broken (endothermic) Hydrogen bonds between the water molecules are broken (endothermic) Ion-dipole attractions form between ions and polar water molecules (exothermic) −mcΔT To measure the ​enthalpy of dissolution​, use: ΔH aqueous solution = n. The molar heat of neutralisation​ ΔH neut is the energy liberated per mole of water produced during a q neutralisation reaction. It is calculated by ΔH neut = n(water produced) The implications of water's specific heat capacity on Earth and its climate; Lakes and oceans are slow to freeze. They must lose a lot of energy for the temperature to fall to freezing point. Oceans and other bodies of water moderate the climate by acting as a thermal reservoir to absorb heat from the air without large changes in the temperature of the water. Cities near the ocean will have less extreme temperatures. It is possible for aquatic life to survive strong climate fluctuations as temperature variations in water are not extreme.  Heat capacity and density means that waterways don’t freeze solid which makes aquatic life viable under ice. The temperature of water in ponds stays the same overnight. Water in a car radiator can absorb a large amount of heat from the engine of the car without the heat rising too much. Since the Earth’s surface is 70% water, oceans buffer atmospheric temperatures, providing a liveable climate. The ​enthalpy of vaporisation​ ( ΔH vap ) is the amount of energy required to transform 1 mol of a substance from a liquid into a gas at 1atm pressure. All values will be given to calculate it. The ​latent heat of fusion​ ( ΔH f us ) is the energy needed to melt 1 mol of a solid to a liquid at 1atm pressure. Data values will be given to calculate it. Bond energy:​ refers to the amount of energy required to break 1 mol of a covalent bond into its constituent gaseous atoms at SLC. Bond energies will be provided. To calculate ΔH using bond energies, add up the energy required to break the reactant bonds and subtract the energy released when making the product bonds. Some bond energies are stronger than others due to a smaller atomic radius. Hess’s Law:​ The enthalpy change for a reaction is independent of the reaction pathway taken. In the following reaction, ΔH 1 = ΔH 2 + ΔH 3 When determining ΔH using Hess’ law, some reactions may need to be reversed. ​When a reaction is reversed, the sign of ΔH changes​. The standard enthalpy of formation​ ( ΔH of ) is the change in enthalpy that occurs when 1 mol of a compound is formed in its standard state from its constituent elements in their standard state at SLC. These values will be given. When a substance is in its elemental state, its enthalpy of formation is 0. To calculate the enthalpy change using enthalpies of formation use the formula: ΔH oreaction = ∑ ΔH of(products) − ∑ ΔH of(reactants) The types of ​systems​ are: Open systems - energy and matter can be exchanged with the surroundings Closed systems - only energy can be exchanged with the surroundings Isolated systems - neither matter nor energy can be exchanged with the surroundings You will need to know how to calculate the enthalpy change for: Photosynthesis (endothermic)  ○ 6CO2(g) + 6H 2 O(l) → C 6 H 12 O6(s) + 6O2(g) ΔH = 2804.8 kJ/mol Respiration (exothermic) ○ C 6 H 12 O6(s) + 6O2(g) → 6CO2(g) + 6H 2 O(l) ΔH = − 2804.8 kJ/mol Laws of Thermodynamics: The First Law of Thermodynamics: Energy cannot be created or destroyed in an isolated system The Second Law of Thermodynamics: The entropy of any isolated system always increases with time The Third Law of Thermodynamics: The entropy of a system approaches a constant value as the temperature approaches absolute zero Entropy​ ( S ) represents the degree of disorder in a system or the energy not available to do work and is given in joules per kelvin (J/K). Factors which lead to a change in entropy are: A larger system (with more particles) has more energy than a smaller system A hotter system (more KE) has more entropy than a cooler system A liquid has more entropy than a solid A gas has more entropy than a liquid The ​change in entropy of a system​ ( ΔS ) can be calculated the same way as enthalpy: ΔS osystem = ∑ ΔS oproducts − ∑ ΔS oreactants For any spontaneous process: ΔS total = ΔS system + ΔS surroundings > 0 Gibbs free energy​ ( ΔG ) represents the maximum amount of work a system can complete at constant T and P and is given in kilojoules per mole (kJ/mol). It is calculated by the: Gibbs-Helmholtz equation: ΔGo = ΔH o − T ΔS o where ΔS o must be converted to kJ and T is in Kelvin.. The three outcomes of gibbs free energy are: ○ ΔGo < 0 - The reaction is spontaneous ○ ΔGo = 0 - The reaction is neither spontaneous or non spontaneous ○ ΔGo > 0 - The reaction is nonspontaneous To ​calculate the temperature at which a reaction becomes spontaneous​ let ΔGo = 0 and solve for T given that you have already found ΔH o and ΔS o. Types of enthalpy: o ΔH rxn Enthalpy change of a given reaction at STP o ΔH comb Enthalpy change of combustion of 1 mole of substance with excess diatomic oxygen at STP ΔH of Enthalpy change when 1 mole of a product is produced from its constituents at STP ΔH fous Enthalpy change required to melt or freeze 1 mole of a substance o ΔH vap Enthalpy change required to vaporise or condense 1 mole of a substance ΔH ao Enthalpy required to convert 1 mole of an element in its standard state to its gaseous state at STP o ΔH L(diss) Enthalpy required to convert 1 mole of a solid lattice to its gaseous state at STP