Chemical Reactions Final Study Guide

Chemical Reactions: -​ Law of conservation of mass: mass in a closed system can neither be created nor destroyed Rules for adding coefficients: 1.​ Write out formulas 2.​ Count the number of atoms of each type on both sides 3.​ Balance the element one at a type by adding coefficients in front 4.​ Double check -​ -​ -​ -​ -​ -​ -​ Combination/ Synthesis reaction: A + B -> AB Decomposition reaction: AB -> A + B Combustion reaction: hydrocarbon + O₂ -> CO₂ + H₂O Single replacement reaction: one element replace another (*be extra careful with H₂O) and replaced the element most similar to itself Double replacement: 2 things replace each other Physical change: does not change the identity of the substance Chemical change: results in a rearrangement of atoms, changing the identity of the substance Signs of a chemical change: 1.​ Production of an order 2.​ Change in temperature 3.​ Change in color 4.​ Production of a gas 5.​ Formation of a solid (precipitate) Conversion of Matter: -​ -​ -​ Mole: unit for amount and is defined as the number of carbon atoms in exactly 12g of carbon-12 Avogadro’s number: 6.022 x1023 Molar mass: mass of 1 mol (g/mol) *used to convert from g -> mol Volume may be effected by: 1.​ Temperature 2.​ Pressure -​ STP: Standard Temperature and Pressure, 1 atm of pressure and 273 K -​ Molar volume: at STP 1 mol of gas occupies a volume of 22.4 L Calculating composition: 1.​ List the masses of each element and the total mass -​ Used precision rounding 2.​ Divide each element’s mass by the total mass and x100% (part/ whole) x100% = percent Stoichiometry: Ways to interpret a balanced chemical equation: 1.​ Moles 2.​ Representative particles (atoms, molecules, formula units) 3.​ Mass (g) 4.​ Volume (L @ STP) -​ Mole ratio: uses coefficients in balanced equations to write ratios or conversion factors Ex: 2H2 + O2 -> 2H2O 1.​ mol H2O -> mol O2 ​ (1 mol O2/ 2 mol H2O) -​ limiting reactant: limits the amount of product that can form Ex: if you needed 2 pieces of bread, 1 slice of ham and 2 slices to tomato to make a sandwich and you have 4 pieces of bread, 2 slices of ham, and 3 slices of tomato; tomato would be the limiting reactant ​ Steps to finding the liminating reactant: 1.​ Balance equation 2.​ For each reactant calculate the mole of one product each react will produce * this needs two set ups but no plan is required * use the given amount (usually grams) then using both the molar mass and molar ratio to find the amount of moles of the product 3.​ State your limiting reactant (reactant with less product) * List the limiting react as your given in all stoichiometry question 4.​ State your excess reagent 5.​ Complete stoichiometry Excess remaining = express reagent initial - excess reagent reacted -​ -​ -​ Theoretical yield: an amount determined by stoichiometry Actual yield: the amount gotten in a lab * usually smaller Percent yield: (actual yield/ theoretical yield) x100% Gases and Their Properties: -​ -​ -​ -​ -​ Compressibility: gases are easily compressed due to the space between molecules Kinetic energy: energy an object has because of its motion Kinetic theory applied to gases: 1. Gases are made of molecules -​ These are tiny spheres with insignificant volume 2. Molecules are in constant motion -​ They have kinetic energy 3. Collisions are perfectly elastic -​ When molecules collide they do not lose kinetic energy 4. Temperature is directly proportional to the average kinetic energy of the molecules -​ Increasing temperature increases the average kinetic energy of the molecules Pressure: force per area Gas pressure: caused by the collisions of the particles against the walls of the container Variables for describing gases -​ -​ -​ -​ -​ -​ -​ -​ Variable: Units: Pressure P kPa or atm Volume V L Temperature T K or Cº Amont n Mol or g Amount: pressure and amount of directly proportional - more n -> more P - less n -> less P Volume: pressure and volume are inversely proportional - more V -> less P - less V -> more P Temperature: pressure and temperature are directly proportional - more T -> more P - less T -> less P Boyle’s law: P1V1=P2V2 Charles's Law: V1/T1= V2/T2 Guy-Lussac’s Law: P1/T1 = P2/T2 Combined Gas Law: P1V1/T1 = P2V2/T2 Ideal gas law: PV = nRT - allows us to calculate with moles instead of having to be @STP but there is no change in condition -​ -​ Ideal gas constant (R): 8.31(L x kPa/ mol x K) Dalton’s law of partial pressures: the total pressure (Pt) of a container with multiple gases equals the sum of the other particle pressures Water and Solutions: -​ -​ Intramolecular forces: forces within a molecules (chemical bonds - ionic, covalent) Intermolecular forces: attractive forces between different molecules Water is highly polar since it is very electronegative since oxygen pulls the shared electron closer to itself and give oxygen a partially negative charge and leaves the hydrogens partially positive -​ Hydrogen bonding: H2O “bonds” weakly with other polar or charged molecules * this allows H2O molecules to be attracted to each other because the ∂- from oxygen attracts the ∂+ from the hydrogen of the other H2O molecules Creating molecular diagrams: 1.​ Create a key 2.​ Need relative sizes (use periodic table) 3.​ Use different shading 4.​ Include charges/ particle charges Properties of H2O: -​ Due to the collective strength, the hydrogen bonds absorb a lot of heat which allow HO to resist changes in temperature -​ Cohesion: water molecules “stick” to each other (Hydrogen bonding between H2O and H2O) -​ Surface tension: water forms strong surfaces and spheres due to cohesion -​ Adhesion: water “sticks” to something different than itself (hydrogen bonding between H2O and another polar/ ionic molecule) -​ -​ -​ -​ -​ Solution: homogenous mixture of solvent + solute Solvent: what your dissolving in ex: water Solute: the substance your dissolving ex: sugar Aqueous solution (aq): homogenous mixtures with H2O as the solvent Homogenous mixture: uniform throughout -​ -​ -​ Heterogenous mixture: non-uniform throughout (unevenly) Suspension: a heterogenous mixture with large particles that slowly settle Colloids: a heterogeneous mixture with medium sized particles ex: blood and milk Dissolution explanation: If we want to increase the rate of dissolution we need to increase the contact between solute and solvent particles Factors affecting the dissolution: 1.​ Stirring/ agitation -​ Dispersal increases contact between solute and solvent particles 2.​ Surface area: volume (of solute) -​ Smaller pieces = more surface area -​ Volume = more contact between solute and solvent 3.​ Temperature -​ Higher temperature = greater average kinetic energy of molecules -​ This causes the particles to move faster causing more contact between solute and solvent * this also increases the total amount of solute that can be dissolved -​ -​ Solubility: the max amount of a substance that will dissolve at a specific temperature Unsaturated: can still dissolve more solute ex: solute disappears -​ Saturated solution: contains the max amount of solute dissolved -​ Supersaturated solution: contains more than it theoretically can dissolve Solubility: increase temperature: creates an increase solubility for solids and liquids but decreases solubility for gases -​ -​ -​ -​ -​ Henry’s Law: solubility of a gas in a liquid is directly proportional to the pressure of the gas above the liquid Miscible: when two liquids can mix/ dissolve Immiscible: liquids that cannot mix/ dissolve Concentration: measure of amount of solute dissolved in an amount of solvent Molarity: mol solute/ L solution M = n/V Thermodynamics: -​ -​ Energy: capacity for doing work or supplying heat Heat (q): form of energy that transfer between two objects due to their difference in temperature Types of heat transfer: 1.​ Conduction: heat transferred by particles colliding 2.​ Convection: heat transferred by circulation 3.​ Radiation: heat transferred through waves -​ -​ -​ -​ Endothermic: absorbed energy (heat) and should have a positive value (∆H>0) since its a reactant Exothermic: released energy (heat) and should have a negative value (∆H<0) since its a product Potential energy: energy that is stored in bonds between atoms Joule: SI unit of heat and energy 1 Calorie = 1 kilocalories = 1000 calories 4.184 J = 1 calorie -​ specific heat (C): the amount of heat needed to raise the temperature of 1g of a substance by 1ºc - units are usually J/gºc q= mC∆T C = q/ m∆T -​ q: heat -​ C: specific heat -​ ∆T: change in temperature - ∆T = Tf -Ti -​ -​ Calorimetry: process of measuring heat into or out of a system during a reaction/ change Hess’s law: allows us to find the overall ∆H of a reaction by calculating the sum of the ∆H values of its steps Hess’s Law steps: 1.​ If needed, flip steps (∆H changes signs) 2.​ If needed, multiply steps by any factor (multiplies ∆H) 3.​ Combine like factors -​ Same side = add together -​ Opposite side = subtract * used precision not sig figs -​ -​ Standard conditions (∆Hº): 1 atm, 25ºc, 1M Standard heats of formation (∆Hºf): ∆H of the formation of one mole of a compound from its element in their standard states - Σ∆Hºf products - Σ∆Hºf reactants -​ First law of thermodynamics: application of the law conversation of energy to heat and thermodynamic processes, energy cannot be created or destroyed, but it can be changed -​ Spontaneous/ thermodynamically favorable: naturally occurring under a given set of conditions - low state of energy (-∆H) and more disordered (+∆S) Nonspontaneous/ not thermodynamically favorable: does not occur under a given set of conditions * if a process is spontaneous in the forward direction, the the reserve process will be nonspontaneous Entropy (S): a measure of the disorder or randomness of a system - the units are J/K or J/(mol)(K) Second law of thermodynamics: the entropy of the universe (different that system) always increases for a spontaneous process -​ -​ -​ ∆S is usually positive when: 1.​ Solid -> liquid -> aqueous -> gas 2.​ The total number of gas molecules increases 3.​ A larger molecule is broken into 2 or more smaller molecules (decomposition) -​ Gibbs free energy (∆G): available energy of energy that we can actually use to work - units are kJ or kJ/mol * this contrasts to enthalpy which represents the total energy of the system ∆G = ∆H - T∆S relationship between ∆G and spontaneity: 1.​ ∆G < 0: spontaneous process in the forward direction; releases energy 2.​ ∆G > 0: nonspontaneous process in the forward direction 3.​ ∆G = 0: system is at equilibrium; no net change occurs ∆H: ∆S: ∆G: Reaction outcome: - + Always - The reaction is spontaneous at all temperatures + - Always + The reaction is nonspontaneous at all temperatures - - - at low temperature + at high temperature The reaction is spontaneous at low temperatures + + + at low temperature - at high temperature The reaction is spontaneous at high temperatures Acids and Bases: Properties of acids: -​ Increase [H+] ([H3O]) -​ Taste sour -​ Electrolytes (conduct electricity) -​ React with metals to form H2 gas -​ Changes colors of indicators to red -​ pH < 7 Properties of bases: -​ Increase [OH-] in water -​ Bitter and slippery -​ Changes colors of indicators to blue -​ Electrolytes -​ pH > 7 Nomenclature Name: Formula: Acid or base? Strength: ammonia NH3 base weak ammonium NH4+ acid weak Acetic / Ethnic acid HC2H3O2 acid weak Acetate C2H3O2- base weak Chloric acid HClO3 acid strong Chlorous acid HClO2 acid weak Chromic acid H2CrO4 acid strong Hydrogen Cyanide HCN acid weak Cyanide CN- base weak Dichromic acid H2Cr2O7 acid strong Hypochlorous acid HClO acid weak Nitric acid HNO3 acid strong Nitrous acid HNO2 acid weak Perchloric acid HClO4 acid strong Permanganic acid HMnO4 acid strong Nomenclature Phosphoric acid H3PO4 acid weak Sulfuric acid H2SO4 acid strong Sulfurous acid H2SO3 acid weak Hydrochloric acid HCl acid strong Hydrobromic acid HBr acid strong Hydroiodic acid HI acid strong Potassium hydroxide KOH base strong Sodium hydroxide NaOH base strong Barium hydroxide Ba(OH2) base strong O2- base strong OH- base strong H- base strong Hydrogen ion H+ acid strong hydronium H3O+ acid strong water? H2O Acid or base neutral? Hydrogen fluoride HF acid weak Fluoride F- base weak Aluminum hydroxide Al(OH)3 base weak Hydroxide -​ -​ -​ Monoprotic: contains one hydrogen atoms ex: HCl Diprotic: contains two hydrogen atoms ex: H2SO4 Triprotic: contains three hydrogen atoms ex: H3PO4 Acids: Bases: Strong: Completely ionize Completely dissociate Weak: Ionize only slightly Produce a small amount of of OH- in water Acid: base: Arrhenius Donates H+ in water Donates OH- in water Brønsted-Lowry Donated a proton (H+) Accepts a proton (H+) Lewis Accepts an e- pair Donated an e- pair -​ pH: measurement of [H+] or [H3O+] Formulas for relating to pH or pOH: pH + pOH = 14 pH = -log[H3O+] pOH = -log[OH-] [H3O+] = 10-pH [OH-] = 10-pOH -​ -​ -​ -​ -​ Conjugate acid: formed when a base gains H+ Conjugate base: formed when a acid loses H+ Conjugate acid-base pair: - Strong acids/bases form weaker conjugates - Weak acids/bases form strong conjugates Amphoteric substances: act as both an acid and base Neutralization reaction: when a base and acid react to form salt and water Kinetics and Equilibrium: -​ -​ ​ -​ -​ Reaction rate: the change in concentration of a reactant or product per unit of time Collision theory: 1.​ Molecules must collide with sufficient activation energy 2.​ Molecules must collide in the correct orientation * this causes successful collisions and therefore successful products activation energy (Ea): the minimum energy required for molecules to react * the greater the activation energy the slower the reaction Transition state (キ): activation complex Factors that affect the reaction rate: 1.​ Temperature -​ Higher temperature = higher reaction rate 2.​ Surface area -​ Greater surface area = greater reaction rate 3.​ Concentration -​ Higher concentration of reactant = higher reaction rates 4.​ Presence of a catalyst * all of these help increase the amount of successful collisions -​ -​ Catalyst: substances that speed up reactions but lowering the activation energy Rate laws: the mathematical relationship between the rate of a chemical reaction and the concentration of reactants - only reactants aA -> cC: rate = k[A]n -​ k: the rate constant * only constant when the temperature is constant -​ n: reaction order (how the given reactant effects the rate) * can only be found using experimental data Most common ration orders: -​ 0 order: the rate is independent of [A] -​ 1 order: the rate is directly proportional to [A] -​ 2 order: the rate is proportional to [A]2 aA + bB -> cC + dD: rate[A]m[B]n -​ Le Châteleir’s principle: when a chemical system at equilibrium is disrupted, the system shifts in a direction that minimizes that distances (tries to counteract) Types of disruption: 1. Concentration change: -​ Increasing [reactant] or decreasing [product] will cause the reaction to proceed towards the products (shift right) -​ Decreases [reactant] or increasing [product] will cause the reaction to proceed towards the reactants (shift left) 2. Volume and pressure change: -​ Increased P: reaction will shift toward the direction with fewer gas molecules -​ Decreased P: reaction will shift toward the direction with more gas molecules * if there is an equal amount of gas molecules there is no shift 3. Temperature change: exothermic reactions: -​ Increasing temperature: reaction shifts toward reactants (K decreases) -​ Decreasing temperature: reaction shift toward products (K increases) ​ Endothermic reactions: -​ Increasing temperature: reaction shift toward products (K increases) -​ Decreasing temperature: reaction shift toward reactants (K decreases) Oxidation-Reduction: -​ Oxidation states: indicates the number of charges the atom would have in a molecule/compound if electrons were transferred completely Rules for assigning oxidation states: 1.​ Elements in the r free state have an oxidation number of 0 2.​ Monatomic ions have oxidation states equal to their charge 3.​ In compounds, Fluorine always has an oxidation state of -1 4.​ The oxidation number for oxygen in most compounds is -2 -​ Except for peroxides (H2O2) 5.​ The oxidation number of hydrogen is +1 except when bonded to metals in binary compounds 6.​ The sum of the oxidation states must be zero for a neutral compound or it must equal the overall charge -​ -​ -​ -​ Oxidation: involves the loss of electrons Reductions: involves the gain of electrons Reducing agent: the reactant species that contains the element that is oxidized - since it causes reduction Oxidizing agent: the reactant species that contains the element that is reduced Rules for balancing redox reactions: 0. Assign the oxidation states 1.​ Separate the reaction in half-reactions 2.​ Balance all atoms except H and O 3.​ Balance O atoms by adding one H2O molecule on the opposite side for each O atom needed 4.​ In acidic solutions: balance H+ by adding H on the opposite side 5.​ Balance the charges by adding electrons (e-’s) to the more positive side 6.​ Multiplying 7.​ reactions by the appropriate factor to get the electron to balance out in both half reactions ​ * this allows the electrons (charges) to cancel out 8.​ Add reactions together, subtracting like terms that appear on opposite sides and summing up like terms that appear on the same side