SCH3U Chem Exam Review

Significant DigitsRules:1. all non-zero digits are significant2. all leading and following zeros are not significant3. all zeros between two other digits are significant4. any zero that follows a digit and is to the right of the decimal point is significantWhen performing multiplication or division, the value with the fewest number of significant digits is the number of significant digits used in the final answerWhen performing addition or subtraction, the value with the fewest number of digits after the decimal place represents the number of decimal places in the final answerScientific Notation: used for numbers greater than 1000 or less than 0.001To use, move the decimal to be placed after the first non-zero digitCount the number of decimal places movedIf the decimal has moved to the left, the exponent is positiveTrends of the Periodic TableTrend 1 Atomic Radius:- measured in picometers (1 X 10-12 m)- distance from the nucleus to the valence orbital- going right and up, the atomic radius decreases- this is because going right, there is a greater force of attraction between the protons and electrons which makes a smaller radius- going up the periodic table, there are less valence orbitals meaning the radius becomes smallerKey Terms:Shielding Effect: the effect of filled inner electron orbitals on the attractiveness of valence electrons to the nucleusEffective Nuclear Charge: the apparent nuclear charge experienced by valence electrons, result of the shielding effectTrend 2 Ionization Energy:- amount of energy required to remove an electron- going right and up on the periodic table will increase the ionization energy- this is because fewer valence shells and more protons will make it require more energy to remove electronsTrend 3 Electron Affinity:- energy absorbed or released when an electron is added to a neutral atom- going right and up the periodic table, the electron affinity will increaseTrend 4 Electronegativity:- indicator of the relative ability of an atom to attract shared electrons- smaller atomic radius usually means greater electronegativity- going right and up the periodic table, electronegativity increasesTypes of BondsIonic:- occurs between a metal and non-metal- two ions bond to each other- change in electronegativity is 1.7 – 3.3Polar Covalent:- occurs when non-metals share electrons- has a positive and negative end- change in electronegativity is 0.5 – 1.7Covalent:- occurs when non-metals share electrons- does not have a positive/negative end- change in electronegativity is 0 – 0.5Properties of Ionic and Covalent BondsIonic:- high melting point- does not conduct electricity as a solid- conducts electricity in a solution- high solubility in water- solid state- brittle, hard consistency as a solidCovalent- low melting point- does not conduct electricity as a solid or in a solution- not soluble in water- non-polar compounds soluble in non-polar compounds- polar compounds soluble in polar compounds- can be solid, liquid, or gas- soft, flexible, waxy consistency as a solidLewis Structures- symbol of element is the center- valence electrons are represented by dots surrounding the symbol of the element- when dealing with ionic compounds, arrows are used to represent the loss/gain of electrons- when dealing with covalent compounds, circles around electrons represent the sharing of electrons- line structures simply use lines to show the sharing of electronsNomenclature/Balancing Equations- will not be a major part of the test, so does not need to be reviewed too muchTypes of Reactions:Synthesis- two elements combine to create a binary compound- universal metal + element  compound- multivalent metal + non-metal  compound- non-metal + non-metal  molecular compound- non-metal oxide + H2O will create an acid – will comprise of a hydrogen cation and a polyatomic anion containing oxygen- e.g. SO2 + H2O  H2SO3(aq) (sulphic acid)- metal oxide + H2O  metal hydroxide (base)- e.g. Li2O + H2O  2LiOH(aq)Decomposition:- compound is broken down into simpler ones or elements- binary compound decomposes into elements- metal nitrate decomposing into metal nitrite + O2- metal carbonate decomposing into a metal oxide + carbon dioxideCombustion- chemical reaction of a substance with oxygen producing one or more oxides, heat, and light- inorganic combustion general formula: X + O2  XO- organic combustion only occurs with hydrocarbons which are compounds composed of only hydrogen and carbon- complete combustion of a hydrocarbon general formula: CxHy + O2  CO2 + H2O- incomplete combustion of a hydrocarbon general formula: CxHy + O2  C(s) + CO + CO2 + H2OSingle Displacement- chemical reaction in which one element in a compound is displaced by another element- always refer to reactivity chart when dealing with single displacement reactions- metal displacing another that is below it on the activity series in an ionic solution- e.g. Cu + 2AgNO3  Cu(NO3)2 + 2Ag- metal ion displacing hydrogen in an acidic solution- e.g. Mg + 2HCl(aq)  MgCl2+ H2- active metal + water  metal hydroxide + H2- e.g. 2Rb + 2H2O  2RbOH + H2- halogens displacing other halogensDouble Displacement- chemical reaction in which positive ions of two ionic compounds exchange places, creating 2 new ionic compounds- neutralization reaction: acid + base  water + “Salt” (ionic compound)- e.g. H2SO4(aq) + Ca(OH)2(aq)  H2O + NaCl- precipitation reaction: reaction with two ionic solutions that creates a precipitate- e.g. 2LiCl(aq) + Pb(NO3)2  2LiNO3 + PbCl2SolubilityEffect of Temperature on Solubility- every unique pure substance has its own unique solubility based on the types of bond present- units used to describe/measure solubility is: mass of solute/100mL of solventSolubility of Solid- trend: solubility of solids increase as temperature increases- energy is required to break apart bonds of solids when dissolved in water- as temperature increases, there is more energy to break these bondsSATP/STP: standard ambient temperature and pressureS = solubleSS = slightly solubleI = insolubleSolubility of Gases- trend: solubility of gases decrease as temperature increases- with more energy, the gas particles escape the solutionSolubility of Liquids- trend: solubility of liquids is not affected by the temperature- solute: liquid with less amount- solvent: liquid with greater amountWater- universal solvent- small size, highly polar nature, and the capacity to form hydrogen bondsmakes water very successful at dissolving solutes- water has a permanent dipole- the negative end is attracted to the positive end, causing a special type of attraction called hydrogen bondingHydrogen Bonding- any substance containing hydrogen and oxygen/fluorine/nitrogen- doesn’t create an actual bond, uses strong intermolecular forces to create a force of attraction- hydrogen bonded compounds are likely to dissolve in waterProperties- water is held together by covalent bonds which re stronger than hydrogen bonds- however, the hydrogen bond is stronger than regular dipole-dipole attractions- this results in higher boiling points because more energy is required to break apart these bonds- hydrogen bonds also result in higher surface tensionVolume/Volume Concentration (V/V) %- volume of solute(mL)/volume of solution(mL) X 100%Mass/Volume Concentration (m/V) %- mass of solute(g)/mass of solution(mL) X 100%Mass/Mass concentration (m/m) %- mass of solute(g)/mass of solution(g) X 100%Very Low Concentrations- parts per million (ppm) = mass of solute/mass of solution X 106- parts per billion (ppb) = mass of solute/mass of solution X 109Molarity- molar [ ]- moles of solute/1L of solutionDilutions- reducing concentration of a solute by adding additional solution to the mixture- standard/stock solution: one where the [ ] is known- c1V1 = c2V2- c1 is the initial [ ]- c2 is the final [ ]- V1 is the initial volume- V2 is the final volumeDouble Displacement Reactions- 2 possible outcomes- Compounds remain as ions and no reaction occurs (NR)- New compounds created that consist of 2 of the following: solid precipitate, gas, or waterNet Ionic Equations- an ionic compound dissociates in water and is broken up into its constituent ions- the above occurs before a double displacement reaction happens- net ionic equation only contains the new product and the constituents that produce this compound- spectator ions are any ions not involved in the creation of the new productAcids- sour taste- no texture- conducts electricity in an aqueous solution- pH less than 7- turns litmus paper red- phenolphthalein is colourless- Acid + Metal  H2(g)- Acid + Carbonate  CO2(g) + H2O(l)Bases- bitter taste- slippery texture- conducts electricity in an aqueous solution- pH greater than 7- turns litmus paper blue- turns phenolphthalein pinkArrhenius Theory- an acid is any substance that will ionize in water to produce H ions- a base is a substance that will dissociate in water to produce OH ions- H ions cannot exist alone, and thus exist attached to H2O, creating hydronium: H3O- Only valid for reactions in waterBronsted-Lowry Theory- acids are substances that have an H ion removed- conjugate base is paired with the acid, and becomes the new base- bases are substances that have an H ion added- conjugate acid is paired with the base and becomes the new acidStrong Acids- will completely dissociateWeak Acids- will only have some of the solution dissociated- indicated in a chemical equation by a double arrowMonoprotic Acid- can only give up 1 H ionDiprotic Acid- can give up 2 H ions (H2SO4, H2CO3)Triprotic Acid- can give up 3 H ions (H3PO4)pH and pOH- pH = -log[H or H3O]- [H or H3O] = 10-pH- pOH = -log[OH]- [OH] = 10-pOH- pH + pOH = 14Neutralization- Acid + Base  Salt + Water- Titrations are done to determine the number of moles when the number of moles of H and OH are equal- Equivalence point: the point when titration is complete (H = OH)- End point: a sudden change occurs during a titration- Equivalence is theoretical and determined by calculations- End is experimental and determined by indicatorsGases Kinetic Molecular Theory: explains the behavior of gases in terms of random motion with no volume and no forces of attraction. UP temperature -> breaks bonds faster -> transforming states  Spaces Attraction    Motion TypeShapeVolumeSolid Very tight Strong    VibrationalFixedFixedLiquid  Further Apart Somewhat    Translational, rotational, vibrationalContainerFixedGasVery far apart Weak     Rotational, translational, vibrationalContainerChanges Pressures SI Unit = 1 Pascals (Pa) or Kilopascals (KPa) Older units: torr = 1 mm of Hg STP  = 760 mm Hg = 760 torr = 1 atm = 100.3 KPa SATP = 250C at 100 KPa Temperature Normally in Degrees Celsius, but may have to convert to Kelvin 1 Kelvin = -273.150C Convert C to K -> + 273.15 Convert K to C -> – 273.15   Gas Laws Boyle’s Law Pressure UP = Volume DOWN P1*V1 = P2*V2 Charles’ Law Temperature UP = Volume UP V1/T1 = V2/T2 Gay-Lussac’s Law; Temperature UP = Pressure UP P1/T1 = P2/T2 Combined Gas Law Accounts for pressure, temperature, and volume changes (P1*V1)/T1 = (P2*V2)/T2 Ideal Gas Law properties Elastic collisions Point mass and no volume No forces of attraction Do not condense to liquid Dalton’s Law of partial pressure A pressure that a gas in a mixture would exert if it was the only gas present in the same volume and same pressure Ptotal=P1+P2+P3…+Pn Molar Volume Volume or amount of space that 1 mole occupies at specific T and P Vm=V/n Vm = molar volume V = Volume (in L) n = # of moles Molar volume of ideal gas at: STP: 22.6L/Mol SATP(SAP): 24.8L/Mol Avogadro’s Law States that ALL ideal gases at the same temperature and pressure contains the same number of molecules of gases n1/V1=n2/V2 Ideal Gas LawPV = nRT Where R is the universal constant of gases 8.314 Temperature always in K Volumes in L Pressure in KPa usually recommended Deriving Ideal gas lawMolar Mass mRT/PV = M (g/mol) Densitym/v = MP/RT (g/L) Molar Concentration C = n/V = P/RT (mol/L) Pressure of Water Vapors Pdry gas= Patm – Pwater vapour Pressure of water vapour obtained with a table that relates it to temperature Gas Stoichiometry Balancing Equation Convert everything to moles (from formulas etc.) Mole to mole ratio Solve for your unknown Must find limiting reactant if necessary Volume to Volume: nothing to do since volume = number of moles STOICHIOMETRY Isotopes and average atomic mass Isotope: element variations with different atomic mass but same atomic number Isotopic Abundance: the relative amount in which each isotope of an element Calculating: given 2 isotopes of an element [B: 10.01u] [B: 11.01u]1) Given average atomic mass: 10.81uIf given percentages, you can find average atomic mass by adding each portion 2) Set variables for unknown values Let x represent % abundance for 10/5 B Let 1-x represent % abundance for 11/5 B Calculate: 10.01x+(1-x)11.01=10.81 10.01x+11.01-11.01x=10.81 -1.00x=0.2 X=0.2 The Mole and Avogadro’s constant Mole: the SI unit that is used to measure how much of a substance. (balanced eqn coefficient are moles) Converting Moles -> number of particles Number of Moles = number of particles / Avogadro’s constant Avogadro’s constant: 6.02214179 x 10^23 Converting Moles to Mass/molar mass Molar Mass: how much grams per mol of an element = atomic mass Number of Moles = mass / Molar Mass Percentage Composition Laws of definite proportions: a law stating that a chemical compound always contains the same proportion of elements by mass Mass percentages (H in H2O) = molar mass of H (1.01) / molar mass of H2O (18.02) = 11.2% Empirical Formula A formula that shows the smallest whole number ratio of elements within a compound Lactic Acid: a chemical name Empirical Formula: CH2O Molecular Formula: C3H6O3 Calculations: when given percentage composition of different elements 1) Find the number of moles for each element (assume % is per 100g for simplicity) 2) Divide all the moles by the smallest number mole to reduce the decimals 3) If not whole numbers yet, multiply by some special factors to reduce decimals1.25 x ¼ = 1 4) Rewrite equation into the empirical formula Molecular Formula A chemical formula that expresses the exact ratio of atoms (whole numbers) for a compound NO REDUCTION REQUIRED, and not suitable for ionic bonds (metal + non metal) Determining Molecular Formula: Molar mass of compound must be given Let X represent a multiplier that relates the empirical formula to molecular formula X = molar mass of actual compound / molar mass of EF Multiply all the elements in the empirical formula by that X multiplier Stoichiometry: study of relationships between reactants and products involved in chemical reactions IOWs, study of how much of reactants are required to produce sufficient products Simply use number of moles in reactions (EF) to become their ratios between each otherAl + 3 CuSO4 -> Al2(SO4)3 + 3 Cu  :  2:3:1:3 ratio To determine amounts, use ratio and plug in numbers correctly Limiting/excess reactants Limiting:Reactants that are not completely used up in a reactionWhen this reactant is used up, the reaction stops Excess Reactants:the reactant that is left over after a reaction is overRegardless of how much excess remains, the reaction will stop once the limiting is used When 2 or more reactants seem limiting, one must pick the reactant that is limiting all the others Mass to mass Stoichiometry Converting mass through stoichiometry then to another mass Process: 1) Balance equation 2) Convert mass to moles 3) perform mole to mole ratio 4) convert new mole back to mass Finding Hydrate Given a hydrate, and when hydrate is evaporated, and new mass is taken, how much of the hydrate was there? Process 1) find mole of non-hydrate portion 2) find mass of the water that evaporated 3) find mole of H2O 4) based on ratio, find out how many mols of water should exist Based on the amount of amount of actual results in ratio with theoretical ratio Find the amount that were theoretical based on regular ratios of stoichiometry (OR: divide by the smallest mole to get rid of as much decimals) Percentage YieldHow much percent of the actual product will happen over the theoretical amountPercentage Yield = Experimental yield/theoretical yield x 100% Factors: Competing ReactionOther reactions will happen at the same time, limiting the reactant produced Reaction RateSlow reactions take longer than time allocated to complete, collected immaturely Purity of reactantContaminants in the reaction are unaccounted for in the initial calculations Laboratory Techniques Improper techniques leads to loss of reactants and reactants Soluble solvent in filtrate instead of filter paper Soluble products washed away in rinsing Instruments used weren’t accounted for When looking for masses Look for ratio of compound to water Compare ratio with amount of solvents used, there’s the grams Nuclear reactionsInvolves changes to the nucleus of the reactantsChanges the number of protons or neutrons 3 Types of nuclear reactions Alpha ParticleDecay (a decay) Loss of 2 protons and 2 neutrons Always creates 4/2 He product Beta Particle Decay(b decay) Loss of an electron-like beta particle 0/1 e always a product, adds one neutron to other product Gamma Decay(Y decay)Releases beta particle as well as gamma energy Conditions of nuclear reactions Sum of mass numbers of product = sum of mass numbers of reactants Sum of atomic numbers of product = sum of atomic numbers reactants Nuclear FissionOccurs when an unstable isotope splits into smaller fragments. Multiple neutrons released along with high energy (radiation) Nuclear FusionNucleus absorbs particles, merging 2 atomic nuclei together

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