Chemistry Notes PDF

Chemistry Notes Module 1: Properties and Structures of Matter 1.2: Properties of matter PHYSICAL: properties that can be measured and observed without the substance undergoing a chemical change. CHEMICAL: relate to the chemical reactions that it is able to undergo when exposed to other substance...

Chemistry Notes Module 1: Properties and Structures of Matter 1.2: Properties of matter PHYSICAL: properties that can be measured and observed without the substance undergoing a chemical change. CHEMICAL: relate to the chemical reactions that it is able to undergo when exposed to other substances, heat or light. Examples: toxicity, reactivity, oxidation states, possible bonds, etc CLASSIFYING MATTER PURE SUBSTANCES: represented by chemical formula. E.g Zn (zinc) H2 (hydrogen) CO2 (carbon dioxide), C6H12O6 (glucose). MIXTURES: cannot be represented by singular chemical formula HOMOGENOUS: the same, matter with uniform composition and properties throughout (solution) single phase. EXAMPLES: Vodka, steel, air, rain HETEROGENEOUS: dissimilar, mixture is not uniform in composition or properties. EXAMPLES: salt in water, concrete, blood SEPARATION BY PHYSICAL PROPERTIES Method Physical property it Explanation of Diagram exploits method Solubility Solubility Choosing a solvent that will dissolve one substance, yet not another, will allow filtration and other methods of separation to be performed Filtration State and size A mixture containing an insoluble solid and a liquid is passed through a filter. The filter only allows certain substances of a particular size through Evaporation Boiling point The mixture is heated to the boiling point of the components with the lowest boiling point, enabling it to evaporate and leave behind the other components Distillation Boiling points of Similar to different liquids evaporation, heated to the boiling point of the liquid with the lowest boiling point, that then evaporates and condenses into a separate beaker Decantation Density, immisicbility Immiscible liquids are  separated in a funnel PROPERTIES OF ELEMENTS CLASSIFICTION PROPERTIES Metals ★ Most solid at room temperature ★ Shiny, lustrous appearance ★ Conductors of heat and energy ★ malleable and ductile Non-metals ★ Represented by all states at room temp ★ Most have dull appearance ★ Insulating ★ Neither malleable or ductile semi-metals ★ Known as mettaloids or semi-metals ★ Some have properties of both metals and non-metals ★ Present in specific pattern on periodic table GROUPS AND PERIODS ★ Groups are the columns, periods are the rows GROUP PROPERTIES Group 1: Alkali metals (first column) ★ Very reactive, only one electron way from full valence shell ★ Reacts with water to form strong bases/alkalis ★ Loses 1 electron to make stable ion of charge +1 Group 2: alkali earth metals (second column) ★ Somewhat reactive, 2 electrons away from full valence shell ★ Reacts with water to form bases/alkalis ★ Loses 2 electrons to form stable ion charge of +2 Group 13: boron group “Triels” (13th column) ★ Lower reactivity than group 2 ★ Do not react with water vigorously, but will oxidise over time ★ Lose 3 electrons to give ion charge of +3 Group 14: carbon group, “Tetrels” (14th ★ Lower reactivity than group 13 column) ★ Do not react with water ★ Group 15: nitrogen group, Pnictogens (15th ★ Higher reactivity than group 14 column) ★ Tend to react with oxygen over time to form stable ions, such as nitrate ★ Gain 3 electrons to have a 3- charge Group 16: Oxygen group, Chalogens ★ Higher reactivity than group 15 ★ Oxides and sulphides tend to react with water and oxygen to form hydrogen sulfates ★ Gain 2 electrons as an ion to form 2- Group 17: Halogens ★ Very reactive ★ Often occur in nature as diatomic molecules ★ Gain 1 electron to become an ion charge of 1- Group 18: Noble gasses ★ Very unreactive ★ Dont form molecules or ions as they have complete valence shell ★ Only exist as singular atoms ★ Always gases Group 3-12: transition metals ★ Properties of metals ★ No large trends ★ Have differing charges as ions ★ Except for zn 2+ or Ag + Group 3-13: Lanthanides and actinides ★ Generally less common ★ Hard to extract and separate from each other ★ Elements beyond 92 are not found  naturally 1.3 Atomic Structure ATOMIC NUMBER Z: number of protons in nucleus MASS NUMBER = number of protons + number of neutrons NUCLEONS: particles that reside in the nucleus (proton or neutron). E.g atom of carbon has 12 nucleons ISOTOPES: atoms of the same element that differ in number of nuetrons, some may be radioactive RELATIVE ATOMIC MASS = isotopic abundance1/100 x mass number1 + isotopic abundance2/100 x mass number2 +... RADIOACTIVITY ★ Caused by an instability in the nucleus due to: - too many neutrons  - too few neutrons - being too large ★ When unstable, radioactive isotopes can be called radioisotopes FORCE STRENGTH DISTANCE DETAILS Strong nuclear force Very strong Very short Attractive, only within nuclei electromagnetism strong Infinite Attractive or repulsive Weak nuclear force weak Short Governs radioactive decay gravity Very, very weak Infinite Attraction between masses ★ The stability of a nucleus is the result of balance between attractive strong force and electromagnetic force When nucleus is small, attractive force is greater than repulsive force, allowing for a stable isotope When nucleus is large, repulsive force is greater than attraction, causing the nucleus to fall apart - unstable, radioactive isotope ★ Unstable nuclei undergo radioactive decay to become stable ALPHA PARTICLE ★ 2 protons + 2 neutrons ★ Helium nucleus BETA PARTICLE β- ★ Electron e- ★ Ejected when neutron decays into a proton GAMMA RAY γ ★ High energy electromagnetic radiation HALF LIVES ★ Very unstable nuclei mean that these radioactive elements have a very short half life ★ More unstable nuclei mean that these radioactive elements have a very long half life SCHRODINGER MODEL ★ Electrons are waves, not particles ★ Electron orbits are quantised ★ Electron orbitals are not definite paths, but are probability clouds ★ Explains why atoms are stable and do not collapse ★ Predicted emission spectra ★ Explained and informed general shape and meaning behind periodic table ORBITAL NOTATION 1.4 Periodicity ★ Across periods, elements typically move from solid to gaseous state due to types of bonds and structure CORE CHARGE: Measure of attractive forces between the valence electrons of an atom and the nucleus Core charge = protons - inner shell electrons Trend: increase across period as number of protons increases, remains the same down a group ATOMIC RADII: Measurement from the centre of the nucleus to the edge of the outermost occupied electron orbital Trend: ★ decreases across a period - core charge increases, number of electron shells remains the same - valence electrons are more strongly attracted to the nucleus ★ increases down a group - core charge is constant, yet number of electron shells increase - valence electrons are less attracted to the nucleus IONISATION ENERGY - The amount of energy required to remove most loosely held electron Trend: ★ Decreases down a group - elements get bigger, more electron shells - valence electrons are further from nucleus, meaning there is a weaker attraction - inner shells sheild attraction ★ Increases across a period - nucleus gets larger, more protons, increase in attraction - extra electrons are of same energy levels, dont experience as much shielding ELECTRONEGATIVITY: A measure of the ability of an atom to attract electrons ★ Increases towards the top right of periodic table ★ Results in unequal sharing, SEPARATION OF CHARGE 1.5 Chemical Bonding INTRAMOLECULAR BONDING ★ Bonds between atoms that form molecules ★ Atoms proceed to an electron configuration of the highest ionisation energy to achieve noble gas configuration for their electrons (full valence shell) ★ Outright transfer of electrons forming ions, leading to IONIC BONDING. This occurs from opposite sides of the periodic table ★ Sharing electrons in larger molecular orbits, COVALENT BONDING. This occurs between non-metals IONIC COMPOUNDS 1. Determine the formula and charges of each ion 2. Determine the ratio they will attract each other to make a neutral compound (zero charge) 3. Write the formula using the numbers for each ion as a subscript number E.g Na2S ★ Monatomic ions are ions of singular atoms, e.g Li+ ★ Monatomic anions end with the suffix ‘-ide’ ★ A polyatomic ion is a group of atoms that are bonded covalently, but have an unequal number of electrons and protons and are therefore charged ions POLYATOMIC ATOMS - NEED TO REMEMBER COVALENT BONDING ★ Share pairs of electrons until they have their valence sells complete ★ If the atoms valence shell is still not shared, each atom will donate another electron to form a double covalent bond, triple bond, etc until full BOND POLARTY ★ When electronegativity of bonding atoms is unequal, bonding electrons are shared unequally ★ Non-polar covalent bonds are the strongest polar bonds as there is maximum electron density between the two atoms. As the electronegativity increases, the bond strength decreases. Ionic bonds are stronger than covalent bonds, they are the strongest bond PROPERTIES OF IONIC AND COVALENT COMPOUNDS ★ Melting an ionic solid needs to overcome the very strong electrostatic attraction between ions, meaning that ionic compounds have very high melting points ★ Ionic substances are hard because of this strong ionic bonding ★ Electrical conductivity occurs in substances that have ions that can move freely. An ionic solid has fixed ions, and will therefore be an electrical insulator ★ As a liquid and dissolved in a solvent, ionic compounds will conduct electricity as their ions are now mobile ★ The properties of covalent compounds are dependant on the attraction forces between molecules - often called intermolecular forces METALLIC BONDING ★ Bonding between metals ★ When metal atoms come together, they form an orderly matrix of nuclei, surrounded by a delocalised mass of valence electrons that can flow and associate with any nucleus ★ This leads to macroscopic properties of metals ★ Electrical conductivity - electrons flow very freely ★ Thermal energy ★ When metal structure is distorted, electrons can move with the distortion and keep structure coherent, which is why metals are malleable and ductile 1.6 Intermolecular forces and allotropy PREDICTING THE SHAPES OF MOLECULES VSEPR - valence shell electron-pair repulsion theory ★ Electrons mostly come in pairs - “lone pairs” and “bonding pairs” ★ The locations that these electron pairs are located in a sphere around an atom will be as far away as possible from each other as electrons repel VSEPR PROCEDURE 1. Draw lewis dot structure of the molecule or polyatomic ion 2. Determine how many bonding groups there are around the central atom and determine molecular basis 3. Determine how many lone pairs are in the shape 4. Use the number of bonding groups vs the number of lone pairs to describe the molecules final geometry POLAR BONDS TO POLAR MOLECULES ★ A polar molecule can also be called a dipole ★ The more electronegative atom is labeled with a partial negative charge, and the more electropositive is labelled with a partial positive charge ★ Draw vectors with the head point at the more electronegative end ★ Once we have assigned any polar bonds and drawn vectors, it can be decided if we have a polar molecule or if the molecule has a ‘dipole moment’ or ‘dipole’ POLAR MOLECULES: both polarity of bonds and shape of molecule ★ Symmetrical molecules are non-polar - bonds are both equal and opposite EXAMPLE: CO2 ★ Asymmetrical molecules are polar - bonds are either not equal or not opposite EXAMPLE: H2O INTERMOLECULAR FORCES ★ The forces between adjacent molecules, atoms or ions ★ Electrostatic attraction and no electrons are shared or transferred ★ Unequal distribution of electrons - like poles repel and opposites attract Main types: ★ Dipole-dipole ★ Dispersion ★ Ion-dipole ION DIPOLE FORCES  ★ Between an ion and a polar molecule ★ Ion has permanent charge, dipole has partial charge (more positive and more negative end) ★ A positively charged ion will have an attraction to the more negative end of a dipole and a repulsion to its positive end, and vice versa for a negatively charged ion ★ Ion dipole forces tend to be the strongest attraction of all intermolecular forces ★ Example of ion-dipole forces is dissolution of ions in water. DIPOLE DIPOLE FORCES ★ Between two polar molecules ★ The more polar molecules involved, the stronger the dipole-dipole attraction between them ★ Very strong dipole-dipole attraction is called HYDROGEN BONDING ★ Very polar molecules tend to have a highly electronegative atom (F,N,O) covalently bonded to a hydrogen atom  ★ FOR EXAMPLE: water ION INDUCED DIPOLE FORCE ★ Between ions and non-polar molecules ★ As ion approaches a non-polar molecule it will cause electrons in the non-polar molecule to move, producing a slightly polar molecule for a moment ★ A cation will attract all the electrons in a non-polar molecule, while an anion will repel the electrons in a non-polar molecule away ★ Generally weaker than dipole-dipole forces DIPOLE INDUCED DIPOLE FORCES ★ Occur between polar and non-polar molecules ★ When polar molecule approaches a non-polar molecule, the electrons in non-polar molecule move producing a slightly polar molecule for a moment ★ Dipole induced dipole forces can be thought of as short term dipole-dipole forces - it only exists when the molecules are close and are generally much weaker than dipole-dipole forces DISPERSION FORCES ★ Between any molecule ★ Small attractive force that arises from a temporary small dipole formed between two molecules that are close together ★ Weakest of all intermolecular forces ★ Two separated non-polar molecules at a distance have very little attraction, however as they move closer, the attractive forces between the nuclei and the opposite molecules electrons starts to become closer, forming two very slightly polar molecules ★ Once molecules are close enough there will be electron-electron repulsion so they will move away from each other ★ The longer the molecules, the greater the dispersion force  ★ The straighter the molecules the greater the dispersion force between them SOLUBILITY ★ Intermolecular forces between molecles govern the physical properties of solubility ★ For a substance to be soluble in another, all intermolecular forces must be of similar strength MELTING POINTS AND BOILING POINTS ★ Stronger intermolecular forces need more energy to overcome ★ Particles that have strong intermolecular forces have higher metling and boiling points ALLOTROPY ★ Allotropes are differnt forms of one element that have distinctly different physical properties ★ EXAMPLE: carbon exists as both graphite and diamond Diamond: covalent network lattice of carbon atoms bonded in a tetrahedral arrangement with single bonds. Graphite: layered, planar structure of hexagonal lattices with dipole-dipole attractions between each layer ★ EXAMPLE 2: Oxygen and ozone Module 2: Introduction to Quantitative Chemistry 2.7 Chemical reactions, stoichiometry and the mole STOICHIOMETRY *QUANTATIVE CHEMISTRY QUESTION: Aluminium bromide reacts with potassium to form potassium bromide and aluminium If I had 5.00g of aluminium bromide, how much potassium would I need to extract all the aluminium? CONSERVATION OF MASS ★ Matter is neither created nor destroyed during a chemical or physical change, it is just rearranged ★ The mass of the reactants equals the mass of the products EQUATIONS ★ Reactants and products do not always combine in 1:1 ratios ★ Basic rules for balancing equations: - add in multiples of the reactant to product molecules untilt he number and identity of the atoms on each side is equal - dont change the identity of any molecules in the reaction or add a new molecule in ★ (s) solid, (l) liquid, (g) gas, (aq) aqueous (dissolved in solution) ★ Stoichiometry tells us the ratios at which different reactions are consumed and different products are produced ★ THESE ARE NOT RATIOS OF MASS *QUANTATIVE CHEMISTRY QUESTION AlBr3(s) + 3K(s) → 3KBr (s) + Al(s) THE MOLE ★ Unit for amount of a substance ★ Differences in mass - number of atoms in 1.00g of hydrogen is different to the number of atoms in 1.00g of uranium due to differences in the mass of atoms ★ The relative mass of carbon is 12.00 ★ Avogadro’s number= 6.02 x 10^23 particles per mole ★ 1.00g of hydrogen contains 6.02x10^23 atoms per mole and 12.0g of 12 carbon contains 6.02x10^23 atoms per mole ★ ONE MOLE OF PARTICLES CONTAINS 6.022 x 10^23 PARTICLES ★ A mole of elements with higher atomic numbers weigh more, yet they contain the same number of atoms ★ The atomic masses on the periodic table have the units grams/mol CALCULATING MOLAR MASS ★ Atoms - read off atomic mass for the element from the periodic table (carbon = 12.01g/mol) ★ Molecules - add together atomic masses (CO2 = 12.01 + 2x16.00) ★ Ions - read off the atomic mass of the equivalent atom ( ★ Formula units - add together atomic masses (NaCl = ) CONVERTING BETWEEN MASS AND MOLES ★ Mass (g) = moles of substance x molar mass ★ Moles (mol) = mass of substance / molar mass ★ EXAMPLE: *QUATNATIVE CHEMISTRY QUESTION Moles of aluminium bromide in 5.00g = 5.00/(26.98 + 3x79.90) = 0.0187 moles CONVERTING MOLES TO NUMBER OF ATOMS OR MOLECULES (PARTICLES) ★ Moles = particles/avogadrio’s number ★ Particles = mol x avogadrio’s number  ★ EXAMPLE: how many atoms are in 1.25g of aluminium? ★ Ions = mol x avogadrios x ions/particle ★ EXAMPLE: how many sodium ions are ther in 0.75g of sodium carbonate? = 0.75/(2x22.99+12.01+3x16.00) x 6.022x10^23 x 2Na+/1NaCO3 = 8.4308 x 10^21 Na ions PERCENTAGE COMPOSITION ★ %A in a compound = (mass of A in one mole of the compound/mass of one mole of the compound) x 100 Example 1: ★ Mass of Fe in 1 mol of Fe2O3 = 2 x 55.85 = 111.7g ★ % iron in Fe2O3 = 111.7x100/159.7 = 69.9% Example 2: ★ % of nitrogen in urea (CH4N2O) ★ Mass of Nitrogen in 1 mol of urea = ★ % nitrogen in urea = DETERMINING FORMULAE ★ Empiracal formula tells us what atoms are present and in waht ratio but may not be the molecular formula ★ Method: - write down masses of elements in given sample - convert mass to moles - divide by the smallest number of moles to get a ratio - if numbers are not close to whole numbers, multiply by a suitable factor so they become approx whole numbers - round off the numbers to give them as integers  STOICHIOMETRY AND MOLES ★ Stoichiometry gives the ratio of moles needed for reactants to moles produced for products ★ One mole of oxygen combines with two moles of carbon to produce two moles of carbon monoxide *QUANTATIVE CHEMISTRY QUESTION Moles of potassium needed to react with 0.0187 moles of aluminium bromide = 0.0187 moles AlBr3 x 3molesK/1moleAlBr = 0.0561 moles of K needed = 2.19g Therfore 5.00g of AlBr3 needs 2.19g of potassium to react completely away to produce aluminium and potassium bromide LIMITING REAGENT ★ The reactant that runs out first is called the limiting reagent, the other reagents are said to be in excess ★ Determining the limiting reagent in a reaction will allow us to determine how much product could be formed ★ Method: - convert smounts of reactants given to moles - choose one reactant to solve “if i have this many moles of reactant, how many moles of the other do I need?” - compare your answer to the problem with the actual value, if you have more than you need it is in excess, and if not it is limitng - IF ASKING HOW MUCH PRODUCT IS FORMED, determine how much product will have formed by using the number of moles of the LIMITING REAFENT to connect by stoichiometry 2.8 Concentration and Molarity SOLUTIONS ★ A solution is a HOMOGENOUS mixture  ★ The solvent ( substance that does the dissolving) forms majority of the mixture, the solute (substance that dissolves) the minority ★ Water is most common solvent ★ A substance dissolves in another beause its intermolecular attraction is of a similar strength to the solvents intermoleular attraction and the attraction between the solvent and solute ★ Concentration - the amount of solute per amount of solution (ratio) ★ More solvent to solute = dilute, more solute to solvent = concentrated ★ MOLARITY AND CONCENTRATION ★ Concentration = solute/solution ★ Molarity = moles/L (same as concentration) ★ Molarity formula: C = n/v Example: 1.32g of KMnO4 was dissolved in water and the volume made accurately to 250.0 mL. Calculate the concentration in g/100mL and g/L Example 2: determine the mass of sodium chloride neede to make a 250.0mL 90.0% (w/v) solution Example 3: calculate the molarity of a solution of 22 g/L NaCl (aq) 22g/L = 0.376mol/L Dilution ★ The act of adding more solvent to a solution ★ The volume will increase, but the moles of solute will remain unchanged ★ The dilution factor is the ratio of dilution volumes ★ Mol/L(conc) x L(conc)= Mol/L (dil) x L(dil) 2.9 Gas Laws Pressure = force/area E.g. 1kg on a pin, large amount of pressure, 1kg on shoe, less pressure Pascal (Pa = N/m^2) More particles = more collision = more pressure Higher temp = higher pressure Lower volume = higher pressure Earths atmosphere exerts pressure on the surface Changes slightly over time, changes with altitude (higher, less particles, pressure drops) Sea level pressure = 101.325 kPa Atmospheric pressure = atm GAY LUSSAC’S LAW (COMBINING VOLUMES) ★ When measured at a constant temperature and pressure, the volume of gases taking part in a chemical reaction show simple, whole number ratios to one another ★ 100ml H2 + 50ml O2 --- 100ml H2O (g) ★ Volume of gases is not conserved during a reaction, but mass is AVAGADRO’S LAW ★ When measured at same temp and pressure, equal volumes of different gases contain the same number of molecules ★ 100ml H2 (2 atoms) + 100ml Cl2 (2 atoms) → 2HCl (4 atoms) Molar Volume of Gas ★ ONE MOLE OF ANY GAS AT SAME TEMP AND PRESSURE OCCUPY SAME VOLUME ★ At 0º C and 100kPa, all gases occupy 22.71 L/mol ★ At 25ºC and 100kPa, all gases occupy 24.79 L/mol Example: what volume will 50.0 grams of oxygen gas occupy at 100.0kPa and 0ºC? Answer: 22.71L/mol x mol/g x g = 22.71L/mol ÷ molar mass x mass = 22.71 ÷ (2 x 16.00)g/mol x 50.0g = 35.5 L (3sf) Example 2: how many moles are contained in 2.400 L of methane (CH4) at 100.0kPa and 25ºC? Answer: Mol = L x mol/L = L ÷ L/mol = 2.400/24.79 = 9.681 x 10^-2 moles BOYLE’S LAW ★ Pressure is inversely proportional to volume ★ P x V = k(constant) ★ P = k/V ★ i.e if pressure is doubled, volume is halved, if pressure decreases to the fifth, volume is five times larger ★ Therefore, P1 x V1 = k, P2 x V2 = k … P1 x V1 = P2 x V2 ★ Pressure vs volume graph is hyperbolic, pressure vs inverse pressure is linear Example: a balloon had a volume of 5.50 L at 4000m altitude (60.0kPa). What volume would the balloon occupy at sea level (100.0kPa) at same temp? Answer: 60.0kPa x 5.50L = 100.0kPa x V2 = 60.0kPa x 5.50L/100.0 = V2 = 3.30 L (3sf) CHARLES’ LAW ★ For a fixed quantity of gas at a constant pressure, volume increases linearly with temperature ★ Vt = kT Volume, k constant, Temp ★ T(KELVIN) = T(C) + 273.15 e.g 25ºC = 298.15K Example: the volume of gas sample at 0.0º C is 2.31 L. What volume would the same sample have at 82ºC? Answer: 2.31/(0.0 + 273.15)K = V2/(82 + 273.15)K V2 = (2.31/273.15K) x 82 + 273.15 K V2 = 3.00L (3sf) COMBINING BOYLE’S AND CHARLES’ LAWS ★ PV/T = k Example: a gas sample has a pressure of 152 kPa, volume of 2.3L and temp of 10.0C. The sample was released into a 90L tank at 18.3ºC. What was the final pressure? Answer: 152x2.3/(10.0 + 273)K = Px90/(18.3 + 273)K 349.6/283K x 291.3K = P90 P = ( “ ) / 90 P = 4.0 kPa (2sf) ★ For fixed sample of gas at constant volume, pressure increase linearly with temp ★ This can be rephrased to show the relationship between volume and number of particles n = number of moles IDEAL GAS LAW *need for assessment ★ Combination of each law ★ PV/nT = constant(R) (universal gas constant) ★ PV = nRT ★ R = 8.314 L.kPa / mol.K (on formula sheet) ★ Ideal gas law ignores intermolecular attractions between gas molecules, and presumes all collisions are elastic (no loss of kinetic energy to heat) ★ Assumes particles take up relatively no space Example: a 1.00 L gas cylinder at 25ºC contain 258 g of oxygen at capacity. What is the max pressure the gas cylinder can withstand? Answer: PV = nRT P x 1.00 = (258g/(2 x 16.00)g/mol) x 8.314 L.kPa/mol.K P = ((258g/(2 x 16.00)g/mol) x 8.314 L.kPa/mol.K)/1.00 P = 19986.68091 kPa = 2.00 x 10^4 kPa DEPTH STUDY - DALTON’S LAW ★ Pressure of gaseous mixtures ★ Partial pressure is pressure exerted by only ONE particular gas in a mixture ★ The total pressure of a mixture of gases is the sume of the partial pressures of its pure components Example: a container holds a mixture of oxygen and hydrogen gas. It contains 6.7 moles of O2 and 3.3 mol of H2. The container’s volume is 300.0L and temp is 273.15K. a) What is the partial pressure of hydrogen? P(H2) = nRT/V = 3.3 x 8.314 x 273.15 / 300.0 = 7494.198… / 300.0 = 25 kPa (2sf) Module 3: Reactive Chemistry 3.1 Types of Chemical Reactions Indicators of chemical change: ★ Gas is evolved ★ Odour is evolved ★ Solid (precipitate) is formed ★ Change in colour ★ Change in temp ★ Release of light or EM SYNTHESIS REACTIONS X+Y→Z E.g 2Cu (s) + O2 (g) → 2CuO (s) Corrosion of copper 2H2(g) + O2 (g) → 2H2O (l) Combustion of hydrogen Predicting products: ★ Highly electromagnetic elements (F, O, CL, I) react with metals to form flourides, metal oxides, meta; chlorides, etc ★ Metals do not react chemically with other metals, homogenous alloys ★ Non-metals and non-metals react to form covalently bonded compounds Reactions of metals 3.12 Galvanic Cells and Standard Electrode potentials Inquiry questions: how is the reactivity of different metals predicted Redox reactions ❖ Electrons are transferred and the ion metal precipirates and the solid metal dissolves ❖ When metal in contact with an ion where the ion has less attraction to electrons than the metal, no reaction occurs ❖ Activity series: ranking of how reactive metals are ❖ “Activity” is a relative measure of how easily a metal is oxidised - how easy a metal will become an ion ❖ Most reactive : weakly holds and attracts electrons ❖ Least reactive: strongly holds and attracts electrons ❖ Standard reduction potential table is a more formalised activity series using half reactions ❖ The more positive the value, the higher affinity for electrons, the more it will want to be in its REDUCED form. ❖ Electromotive force (EMF) is given in the units of Volts (V) ❖ For a reaction to occur, the reduction hald equation must be more positive than the oxidation ❖ To qualify if we get a spontaneous equation we can calculate the difference in the potentials of our two half equations, if total cell potential is greater than 0V Galvanic Cells - Redox reaction involves the transfer of electrons between two chemical species - In a solution, this occurs by ions being able to move freely through the solvent - Electrons are carried from one species to the other through ions - Galvanic cells separate the oxidation and reduction halves  ❖ Connects redox reactions via a metal wire in the aim of converting chemical energy to electrical energy ❖ Two separate solutions containing electrolytes (substance that in solution will transmit an electrical current, contains ions that can move) have an electrode sitting in them, a wire connects both electrodes, a salt bridge spans the distance between the two solutions to form an ELECTRICAL CIRCUIT ❖ The negative end of the cell is called an anode - electrons are a product ❖ Electrons move towards the positive end, the cathode - electrons are a reactant ❖ The anode material may take part in the reaction where it dissolves, the cathode cannot be chemically changes as metals cannot be reduced, most commonly platinum, carbon or gold ❖ A salt bridge provides electrical contact between the two separated reaction vessels ❖ Working galvanic cells are spontaneous ❖ Galvanic cell can also be made in a u tube so that solutions intermingle, this means salt bridge is not needed Cell notation: X I X^+ II Y^+ I Y Cu I Cu^2+ II Ag^+ I Ag Module 4: Drivers of reaction 4.14 Energy changes in chemical reactions Inquiry question: what energy changes occur in chemical reactions System and surroundings System - the particles that take part in the chemical or physical change Surroundings - the environment around it ❖ Sign of chemical reaction is temperaure change in surroundings ❖ Exothermic: detected when surroundings increase in temperature