Currituck County Schools / Overview



|Chapter 1 Summary Notes |

|Main Concepts |Explanations |

| |-Law of Constant Composition: Ex. In pure H2O, H and O combine in a 1:8 mass ratio. Does law of |

|Elements: substances that cannot be decomposed into simpler |constant composition hold good for CuSO4.5H2O? Why or why not? |

|substances |-Law of Multiple Proportions: Explain the following in terms of multiple proportions: |

|Compounds: substances composed of two or more elements |[pic] |

|Law of Constant Composition or law of definite proportions: the| |

|relative masses of elements are fixed in a given chemical |-Separation Techniques: |

|substance. |Hand Separation- for mixtures that can be visually differentiated based on mass, color, shape etc. |

|Law of Multiple Proportions: Applies ONLY when two elements |Filtration: Fitrate, precipitate, heterogeneous mixtures |

|combine to form two or more compounds. The masses of one |[pic][pic][pic][pic][pic] |

|element which combine with a fixed mass of the second element | |

|are in a ratio of whole numbers |Separating Funnel: For immiscible liquids, layers separate |

|Mixtures: combinations of two or more substances |with lesser density layer on top. |

|Techniques for separating mixtures: filtration, distillation, |Centrifugation: Separates particles of different masses based |

|chromatography |on centrifugal force. Heavier particles at the bottom and the |

|Properties: |lighter particles on top. |

|Physical vs. chemical: Did the sample (really) change? |Distillation- uses differences in the boiling points to separate a |

|Intensive vs. extensive: Does the measurement depend on |homogeneous mixture. |

|quantity of sample? |Chromatography-separates homogenous mixtures (mostly inks) |

| |based on the differences in solubility of the mixture in a |

| |solvent. There is a stationary and a mobile phase |

| |[pic][pic] |

|Chapter 2 Summary Notes |

|Main Concepts |Explanations |

|Naming Compounds Review: Before naming a compound, it is |[pic] |

|important to know its type because naming depends upon the |Practice: 1. Sodium Sulfide (Na2S), Potassium Nitrate (KNO3), Ferrous Sulfate Fe (SO4), Ammonium |

|type. For naming purposes, we classify compounds as ionic |Chloride (NH4Cl), Phosphoric Acid (H3PO4) |

|compounds, molecular compounds, and acids. |2. What is the O.N. of P in PO43- ion? |

|Ionic Compounds can be identified by the presence of a metal in|3. What is the O.N. of Fe in Fe(NO3)3? |

|it. (generally solids) Ex. NaCl, K2SO4, PbSO4 | |

|Molecular compounds are made up of all non metals. (generally | |

|liquids and gases) Ex. H2O, N2O5 | |

|Acids begin with H (generally present as aq solutions or gases)| |

|Ex. HCl, H2SO4, HClO3 | |

|Coordination compound: compound that contains a complex ion or | |

|ions. | |

|Ex. [Cu(NH3)4]Cl2 | |

|Name cation before anion; one or both may be a complex. | |

|(Follow standard nomenclature for non-complexes.) | |

|Within each complex (neutral or ion), name all ligands before | |

|the metal. | |

|-Name ligands in alphabetical order | |

|-If more than one of the same ligand is present, use a | |

|numerical prefix: di, tri, tetra, penta, hexa, … | |

|-Ignore numerical prefixes when alphabetizing. | |

|In any (uncharged) atom: The #f protons= atomic number (Z) | |

|# of e= # of p | |

|Mass # (A)= atomic #- # of neutrons | |

|Atomic Symbol: 6C12 | |

|Isotopes are atoms of the same element containing different | |

|numbers of neutrons and therefore having different masses. | |

|Chapter 3 Summary Notes |

|Main Concepts |Explanations |

|Atomic mass units: | |

|1 amu = 1.66054 x 10-24 g | |

|or 1 g = 6.02214 x 1023 amu | |

|1 atom of 12C isotope defined as weighing exactly 12 amu | |

|[pic] | |

|Percent Composition |Ex. 12C: 98.892% x 12 amu =11.867 |

|Moles: Avogadro’s Number = 6.022 x 1023 = atoms in exactly 12.000 g of 12C = 1 mol |13C: 1.108% x 13.00335 amu =+ 0.1441 |

|-Moles important because they can be used for ratios. Mass cannot be used for ratios. Converting: |AW = 12.011 amu |

|grams → moles → molecules |Therefore 12.011 g C = 1 mol C |

|Determining Empirical Formula | |

|Start with the number of grams of each element, given in the problem.  | |

|-If percentages are given, assume that the total mass is 100 grams so that the mass of each element|Ex. What is the percent of O in CuSO4.5H2O? 57.7% |

|= the percent given. | |

|-Convert the mass of each element to moles using the molar mass |Ex. How many molecules of H2O in 100.0 g H2O? 3.343 x 1024 |

|-Divide each mole value by the smallest number of moles calculated.  |molecules H2O |

|-Round to the nearest whole number.  This is the mole ratio of the elements and is represented by | |

|subscripts in the empirical formula.  | |

|-If the number is too far to round (x.1 ~ x.9), then multiply each solution by the same factor to |Ex. What is the empirical formula of a compound that is composed |

|get the lowest whole number multiple. If one solution is 1.5, then multiply each solution in the |of 80.% Carbon and 20.% Hydrogen? If the molar mass is found to |

|problem by 2 to get 3.If one solution is 1.25, then multiply each solution in the problem by 4 to |be 30 g/mol, what is the molecular formula? CH3, C2H6 |

|get 5.  | |

|Chapter 3 Summary Notes Contd. |

|Main Concepts |Explanations |

|EF from Combustion Data: To find out the EF of a compound, the compound is burned in the |Ex. A 0.6349 g sample of the unknown produced 1.603 g of CO2 and 0.2810|

|air. The equation for combustion of these compounds is all follows. |g of H2O. Determine the empirical formula of the compound. Ans. C7H6O2 |

|For Hydrocarbons: CxHy + O2( CO2 + H2O |Ex. When carbon-containing compounds are burned in a limited amount of |

|For Compounds Containing CHN: |air, some CO(g) is produced as well as CO2 (g). A gaseous product |

|CxHyNz+ O2( CO2 + H2O+ NO2 |mixture is 35.0 mass % CO and 65.0 mass % CO2. What is the mass % C in |

|For Compounds containing C, H and O |the mixture? 32.7% C |

|CxHyOz + O2( CO2 + H2O |Ex. If 6.0 g hydrogen gas reacts with 40.0 g oxygen gas, what mass of |

|[pic]OR |water will be produced? |

|[pic] |2H2 (g) + O2 (g) → 2 H2O (g) |

|Limiting Reactant: The reactant that runs out first. LR can be determined by comparing the | |

|mole ratios of the reactants. LR determines the amount of the products formed. |Ex. A mixture of morphine (C17H19NO3) and an inert solid is analyzed by |

|[pic] |combustion with O2. The unbalanced equation for the reaction of morphine|

|[pic] |with O2 is |

|[pic] |C17H19NO3 + O2 → CO2 + H2O + NO2 |

|Determining the Formula of a Hydrate: To determine the formula of a hydrate, a certain mass |The inert solid does not react with O2. If 4.000 g of the mixture yields|

|of hydrate is heated to drive off the water. Then the mass of water driven-off is |8.72 g of CO2, calculate the percent morphine by mass in the mixture. |

|calculated. Now the mole ratio of water to the compound is calculated. |83.1% |

| | |

| |Ex. If in the previous example, only 40.0 g water were formed, what is |

| |the percent yield and percent error? |

| |Ex. When 2.000 gram of Na2CO3.xH2O was heated, 0.914 gram of anhydrous |

| |residue remained. What is the formula of this compound? Ans: |

| |Na2CO3.7H2O   Sodium carbonate heptahydrate [pic] |

|Chapter 4 Summary Notes |

|Main Concepts |Explanations |

|Reactions: To be able to successfully write reactions, you will need to know the following: Solubility Rules , |Synthesis Reactions: |

|Nomenclature Types of Reactions (explained later in this worksheet), How to write net ionic equations MUST KNOW |[pic] |

|For AP Chemistry reaction prediction: |[pic] |

|Always write balanced net ionic equations (meaning dissociate soluble compounds (based on solubility rules), |Decomposition |

|Metal are insoluble and are atomic, written as (s) in these equations. Ex. Mg(s) |[pic]Single Replacement |

|Molecular compounds such as gases (CO2, H2S Etc.) are written as (g) and will not dissociate into ions. |[pic] |

|Water is written as (l) and does not dissociate. |[pic] |

|Ionic compounds may or may not dissociate depending on solubility rules. Ex. PbSO4 insoluble and NaNO3 soluble. |Combustion Reactions: |

|Even a soluble ionic compound may NOT dissociate if it is in solid form (meaning no water present to actually |[pic] |

|dissociate the ions.) | |

|Weak acids and bases partly dissociate or ionize and are written with a reversible arrow. | |

|Remember PSHOFBrINCl. Phosphorus occurs as P4, Sulfer as S8 and rest as diatomic. | |

|While we are reviewing, remember the difference between Zn and Zn2+ and Cl2 and 2 Cl- | |

|Strong acids (HCl, HBr, HI, HNO3, H2 SO4 (first dissociation only!), HClO4 and HClO3) and strong bases (Group 1 | |

|alkali metal hydroxide and Ca, Ba, Sr hydroxides from group 2) dissociate in aq. Solutions. Weak acids and bases | |

|are not dissociated in net ionic equations. |Precipitation |

|Solubility Rules Always soluble: alkalies, NH4+, NO3-, C2H3O2- |[pic] |

|Types of Reactions: Double displacement. Precipitation, neutralization, gas forming. H2CO3 in water = H2O & CO2 | |

|Single displacement or redox replacement: (metals displace metals and nonmetals displace nonmetals) | |

|Combination or synthesis = two reactants result in a single product |[pic] |

|• Metal oxide + water ( metallic hydroxide (base) | |

|• Nonmetal oxide + water ( nonbinary acid |Ex. How many mL of a 3M NaOH solution are required |

|• Metal oxide + nonmetal oxide ( salt |to completely neutralize 20.0 mL of 1.5M H2SO4? |

|Decomposition = one reactant becomes several products |(Start by writing a balanced equation!) Ans. 20.0 |

|• Metallic hydroxide ( metal oxide + water |mL |

|• Acid( nonmetal oxide + water |Ex. How many g of NaOH is required to completely |

|• Salt ( metal oxide + nonmetal oxide |react with 100. mL of 1M HCl? |

|• Metallic chlorates ( metallic chlorides + oxygen | |

|• Electrolysis decompose compound into elements (water in dilute acids or solutions of dilute acids) |[pic] |

|• Hydrogen peroxide ( water + oxygen | |

|• Metallic carbonates --> metal oxides + carbon dioxide | |

|• Ammonium carbonate ( ammonia, water and carbon dioxide. | |

|Hydrolysis = compound reacting with water. | |

|• Watch for soluble salts that contain anions of weak acid the anion is a conjugate base and cations of weak | |

|bases that are conjugate acids. | |

|Reactions of coordinate compounds and complex | |

|• Complex formation by adding excess source of ligand to transitional metal of highly charged metal ion such | |

|as Al3+ Al =4 ligands others 2X ox # | |

|• Breakup of complex by adding an acid ( metal ion and the species formed when hydrogen from the acid | |

|reacts with the ligand | |

|Redox = change in oxidation state= a reaction between an oxidizer and a reducer. | |

|1. Familiarization with important oxidizers and reducers | |

|2. “added acid” or “acidified” | |

|3. an oxidizer reacts with a reducer of the same element to produce the element at intermediate oxidation state | |

|Molarity (M) = moles solute = mol | |

|volume of solution L | |

|Titration is a method to determine the molarity of unknown acid or base. In titration, an acid or base of unknown| |

|molarity is titrated against a standard solution (whose M is known) of acid or base.The end point in a titration | |

|is indicated by a color change by the indicator. Indicators are weak acids or bases and are added in small | |

|quantity (1-3 drops) to indicate the end point. At equivalence point (which should be close to end point), | |

|moles of H+= moles of OH- | |

|M1V1= M2 V2 (sometimes used to get moles , M= moles/L , so moles= M XV) | |

|-What other ways can you get the moles- for a solid acid or base? For a gas? | |

|Electrolyte: substance which, in aqueous solution, ionizes and thus conducts electricity. Ex: salt in water. | |

|Non-electrolyte: substance which, in aqueous solution, does not dissociate and thus does not conduct electricity | |

|Strong & weak electrolytes: conductivity depends on degree of dissociation and equilibrium position: | |

|HA (aq) ↔ H+ (aq) + A- (aq) | |

|Strong = nearly completely dissociated | |

|Weak = partially dissociated | |

|Molecular equation: shows complete chemical equation with states of matter, undissociated | |

|BaCl2 (aq) + Na2SO4 (aq) → 2 NaCl (aq) + BaSO4 (s) | |

|Complete ionic equation: shows complete chemical equation with states of matter, dissociated if appropriate | |

|Ba2+(aq) + 2 Cl-(aq) + 2 Na+(aq) + SO42-(aq) → | |

|2 Cl-(aq) + 2 Na+(aq) + BaSO4 (s)Spectator ions: present in reaction but do not “participate”; depend on | |

|solubility rules | |

|Cl- (aq) and Na+ (aq) | |

|3. Net ionic equation: shows chemical equation without spectator ions | |

|Ba2+ (aq) + SO42- (aq) → BaSO4 (s) | |

|Chapter 5 Summary Notes |

|Main Concepts |Explanations |

|0th and 1st Laws of Thermodynamics: | |

|2 systems are in thermal equilibrium when they are at the same T | |

|Energy can be neither created nor destroyed, or, energy is conserved | |

|Internal Energy | |

|Includes translational, rotational, vibrational energy | |

|Change (ΔE = Efinal - Einitial) is often measured | |

|ΔE > 0: Energy of system increases (gained from surr.) | |

|ΔE < 0: Energy of system decreases (lost to surr.) | |

|ΔE = q + w |Ex. Octane and Oxygen gases combust within a closed cylinder in |

|q = heat added/liberated from system |an engine. The cylinder gives off 1150J of heat and a piston is |

|q > 0 : heat added to system |pushed down by 480J during the reaction. What is the change in |

|q < 0: heat removed from system |internal energy of the system? (Ans: ΔE = -1630J) |

|w = work done on or by the system | |

|w > 0 : work done to system | |

|w < 0 : system does work on surr. | |

|Calorimetry: Measurement of heat flow, experimental technique used to measure the heat transferred | |

|in a physical or chemical process | |

|Calorimeter: the apparatus used in this procedure; two types: constant pressure (coffee cup) and |Ex. How much energy is required to heat 40.0 g of iron (c = |

|constant volume (bomb calorimeter) |0.45J/gK) from 0.0oC to 100.0oC? (Ans: q = 1800J) |

|Coffee Cup Calorimeter: system in this case is the “contents” of the calorimeter and the | |

|surroundings are cup and the immediate surroundings |Ex. 0.500g of Mg chips are placed in a coffee-cup calorimeter and|

|qrxn + q solution = 0 |100.0mL of 1.00M HCl is added to it. The reaction is: |

|qrxn: heat gained/ lost in the chemical reaction |Mg(s) + 2HCl(aq) ( H2(g) + MgCl2(aq) |

|qsolution: the heat gained/lost by solution |The temp. of the solution increases from 22.2C to 44.8C. What’s |

|[pic] |the enthalpy change for the reaction, per mole of Mg? Assume |

| |specific heat capacity of solution is 4.20J/gK and density of the|

|Heat Capacity, C: Amount of heat required to raise T of an object by 1 K |HCl solution is 1.00 g/mL. (Ans: ΔH= -4.64*105 J/molMg) |

|q = CΔT | |

|Specific Heat (or Specific Heat Capacity), c: heat capacity of 1 g of substance | |

|q = mcΔT | |

|Chapter 5 Summary Notes Contd. |

|Main Concepts |Explanations |

|Enthalpy, H: change in heat content of a reaction at constant P | |

|H = E + PV ( ΔH = ΔE + PΔV ( ΔH = (qp+w) + (-w) ( ΔH = qp |Ex. What is the ΔH of combustion of 100g CH4 if ΔHrxno = -890kJ? (Ans: |

|qp = heat content |-5550kJ) |

|ΔH > 0: heat gained from surr. + ΔH in endothermic reaction | |

|ΔH < 0: heat released to surr. + ΔH in exothermic reaction | |

| | |

|Enthalpy of Reaction, ΔHrxn: heat of reaction, extensive property, depends on states of |Ex. What is ΔHrxno of the combustion of propane? |

|reactions and products |C3H8(g) + 5O2(g) ( 3CO2(g) + 4H2O(l) |

|ΔHrxn = -ΔHreverse rxn |Givens: |

| |3C(s) + 4H2(g) ( C3H8(g) ΔH1 = -103.85kJ |

|Hess’s Law: If a rxn is carried out in a series of steps: |C(s) + O2(g) ( CO2(g) ΔH2 = -393.5kJ |

|ΔHrxn = Σ(ΔHsteps) = ΔH1 + ΔH2 + ΔH3 + · · · |H2(g) + ½O2(g) ( H2O(l) ΔH3 = -285.8kJ |

| |(Ans: ΔHrxno = -2219.8kJ) |

|Enthalpy of Formation, ΔHf: heat needed to form substance from its elements | |

|Standard Enthalpy of Formation, ΔHfo: forms 1 mole of compound from its elements in their |Ex. What is ΔHrxno for the combustion of liquid benzene? |

|standard state (at 298K) |C6H6(l) + 15/2O2(g) ( 6CO2(g) + 3H2O(l) |

|ΔHf of pure element (C, O2, H2, etc) is 0. |Givens: |

| |ΔHfo(C6H6(l)) = +49 kJ/mol |

|ΔHrxno = Σ [n*ΔHfo(products)] - Σ[n*ΔHfo(reactants)] |ΔHfo(CO2(g)) = -394 kJ/mol |

| |ΔHfo(H2O(l)) = -286 kJ/mol |

|Bond Enthalpy: Amount of energy required to break a particular bond between two elements in |(Ans: ΔHrxno = -3268 kJ/mol) |

|gaseous state | |

|Indicates the “strength” of a bond |Ex. What is ΔHrxno for the following reaction? |

|ΔHrxn ≈ Σ [ΔHbonds broken] - Σ[ΔHbonds formed] |CH4(g) + Cl2(g) ( CH3Cl(g) + HCl(g) |

|NOTE: this is the “opposite” of Hess’s Law |Given: |

| |ΔH/mol of C-H bond: +413 kJ/mol |

| |ΔH/mol of H-Cl bond: +431 kJ/mol |

| |ΔH/mol of C-C bond: +348 kJ/mol |

| |ΔH/mol of Cl-Cl bond: +242 kJ/mol |

| |ΔH/mol of C-Cl bond: +328 kJ/mol |

| |ΔH/mol of C=C bond: +614 kJ/mol |

| |(Ans: ΔHrxno ≈ -104 kJ/mol) |

|Chapter 6 Summary Notes |

|Main Concepts |Explanations |

|Electromagnetic Spectrum: radiant energy can travel without matter | |

|[pic]λν c = speed of light = 3.0 x 108 m/s | |

|λ = wavelength (m) | |

|ν = frequency (Hz) | |

|Planck’s Theory: Blackbody radiation can be explained if energy can be released or absorbed in | |

|packets of a standard size called quanta | |

|h = Planck’s constant = 6.63 x 10-34 J-s | |

|Photoelectric Effect: As first explained by Einstein in 1905, the photoelectric effect is the | |

|spontaneous emission of an electron from metal struck by light if the energy is sufficient | |

|Atomic Emission Spectra: spectrum for specific wavelengths of light emitted from pure substances | |

|Bohr’s Model of the H Atom: Bohr applied idea of quantization of energy transfer to atomic model, | |

|theorizing that electrons travel in certain “orbits” around the nucleus | |

|Allowed orbital energies are defined by: |Atomic Emission Spectra for H[pic] |

|[pic] n = principal quantum number = 1, 2, 3 ... |Line Series |

|RH = Rydberg’s constant = 2.178 x 10-18 J |Transition down to (emitted) |

|Line Series: transitions from one level to another |or up from (absorbed)… |

|Heisenberg’s uncertainty principle: The position and momentum of a particle cannot be |Type of EMR |

|simultaneously measured with accuracy. | |

|Schrödinger’s wave function: Relates probability ([pic]2) of predicting position of e- to its |Lyman |

|energy. |1 |

|[pic] |UV |

|Matter as a Wave: | |

|m = h / cλ |Balmer |

|Particles (with mass) have an associated wavelength |2 |

| |Visible |

|λ ’ h / mc | |

|Waves (with a wavelength) have an associated mass |Paschen |

|and velocity |3 |

| |IR |

| | |

| |Brackett |

| |4 |

| |Far IR |

| | |

| | |

| |Probability Plots for 1s, 2s, and 3s Orbitals |

| |[pic] |

|Pauli Exclusion Principle: no two charges in an atom can have the same set of four quantum| |

|numbers n, l, m1, ms. |[pic] |

|Effective Nuclear Charge: the net positive charge acting on the outermost electron. |Ex: Give ground state electron configurations for the following: Ni2+ and |

|Shielding Effect: inner electrons shielding the outer electron from the full charge of the|Ni3+ |

|nucleus. |Ans: Ni2+ = [Ar]3d8, Ni3+= [Ar]3d7 |

|Electron Configuration: the way the electrons are distributed among the various orbitals |[pic]Ex: Is Cupric ion dia or paramagnetic? Why? Paramagnetic, due to an |

|of an atom. |unpaired d e. |

|The most stable, or ground, electron configuration is one in which the electrons are in |Mass Spectrograph for Chlorine (Cl2) |

|the lowest possible energy states. | |

|Hund’s Rule: for degenerate orbitals (orbitals with the same energy), the lowest energy is|[pic] |

|attained when the number of electrons with the same spin is maximized. |(chemguide.co.uk) |

|The periodic table is your best guide to the order in which orbitals are filled. | |

|s-block and p-block contain the representative (main group) elements. | |

|The ten columns in the middle that contain transition metals, elements in which d-orbitals| |

|are being filled. | |

|f-block metals are the ones in which the f-orbitals are being filled. | |

|Diamagnetic: paired electrons | |

|Paramagnetic: unpaired electrons | |

|Mass Spectroscopy: Helps identify # and abundance of isotopes and structures of different | |

|compounds. Chlorine has two isotopes, 35Cl and 37Cl, in the approximate ratio of 3 atoms | |

|of 35Cl to 1 atom of 37Cl. You might suppose that the mass spectrum would look like this | |

|but that is not the case because Chlorine consists molecules that fragment [pic] | |

|(chemguide.co.uk) | |

|You could have the following mass fragments--35 + 35 = 70, 35 + 37 = 72, 37 + 37 = 74. So | |

|the actual mass spectrograph will look like the one on the right. | |

|Photoemission Spectroscopy (PES) |[pic] |

|In a photoelectron spectroscopy experiment any electron can be ionized when the atom is | |

|excited. Unlike the first ionization, in this experiment any electron can be removed, not | |

|just the electron that requires the least amount of energy. PES gives insight into the | |

|structure of atom. Each peak in PES indicates the number of electrons and the position of | |

|the peak indicates the amount of energy required to remove those electrons. Note that s | |

|electrons will require more energy than p electrons due to higher ENC hence s electrons | |

|will be farther out on the energy axis. | |

| | |

|Chapter 7 Summary Notes |

|Main Concepts |Explanations |

|Periodic Trends Key Words |[pic] |

|Principal Energy Level: more energy levels means bigger atoms | |

|Nuclear Charge (# of p in an atom): attraction of electrons to | |

|nucleus; increased nuclear charge causes atomic radius to |Atomic Radius |

|decrease |[pic] |

|Shielding Effect: inner electrons increase atomic size by |Isoelectronic Series |

|reducing the attractive force on outermost electrons |Ionization Energy |

|Effective Nuclear Charge: force of attraction felt valence e- | |

|from nucleus; a high ENC means smaller ionic radius (greater | |

|attraction to outermost electrons) | |

| | |

|Atomic Radius: half distance between two covalently-bonded | |

|atoms | |

| | |

|Periodic Trends | |

|Atomic Radius: increase going down, decrease going right (more | |

|energy levels and shielding going down, higher ENC going right)| |

|Size of Ions: cations smaller than anions (fewer e-, less e- | |

|repulsion), bigger going down,smaller going right | |

|Isoelectronic series have same # of electrons, different # of | |

|protons | |

|Ionization Energy: the amount of energy required to remove an | |

|electron from the ground state of a gaseous atom or ion. | |

|A(g) ( A+ + e- | |

|Bigger going right, smaller going down | |

|Exceptions: between Group 2&13, Group 5&6 | |

|Chapter 7 Summary Notes |

|Main Concepts |Explanations |

|Periodic Trends (cont’d) | |

|Electron Affinity: amount of energy required to add an electron|Electron Affinity |

|to a gaseous atom: |[pic] |

|Cl(g) + e− ( Cl− |Electronegativity |

|Larger going right, smaller going down | |

|Exceptions: between Group 1&2, Group 14&15 | |

|Electronegativity: tendency to attract electrons in a covalent | |

|bond | |

|Increases going up and right | |

| | |

|Metals and Nonmetals | |

|Metals: tend to form cations, metal oxides are basic | |

|Nonmetals: tend to form anions, nonmetal oxides are acidic, | |

|poor conductors of electricity | |

|Metallic character increases down a group, decreases across a | |

|period | |

| | |

|Alkali metals (1A)—The most reactive metal family, these must | |

|be stored under oil because they react violently with water! | |

|They dissolve and create an alkaline, or basic, solution, hence| |

|their name. | |

|Alkaline earth metals (2A)—These also are reactive metals, but | |

|they don’t explode in water; pastes of these are used in | |

|batteries. | |

|Halogens (7A)—Known as the “salt formers,” they are used in | |

|modern lighting and always exist as diatomic molecules in their| |

|elemental form. | |

|Noble gases (8A)—Known for their extremely slow reactivity, | |

|these were once thought to never react; neon, one of the noble | |

|gases, is used to make bright signs. | |

| |

|4section6.rhtml | |

|Chapter 8 Summary Notes |

|Main Concepts |Explanations |

|Intramolecular Bonding | |

|Ionic: electrostatic attrction between oppositely charged ions. Generally solids. Ex: NaCl, K2SO4 |[pic] |

|Covalent: sharing of e- between two atoms (typically nonmetals). Generally gases/liquids. Ex: CO2, | |

|SO2 | |

|Metallic: “sea of e-”; bonding e- relatively free to move throughout 3D structure. Generally solid.| |

|Ex: Fe, Al |Ex. Draw the Lewis symbol for fluoride. |

|Covalent Network: large number of atoms/molecules bonded in network through covalent bonding. Ex: |[pic] |

|SiO2, Si, Ge, Diamond, Graphite | |

|Ionic Bonding | |

|Results as atoms lose or gain electrons to achieve a nobel gas electron configuration. Typically | |

|exothermic. | |

|Bonded state lower in energy (more stable) |Ex. In the following pairs, which has a greater lattice energy |

|Opposite charges create electrostatic attraction, which determines the strength of the ionic bond |and why? |

|Occurs when difference in electronegativy is > 1.7 |NaCl or KCl |

|Use brackets when writing Lewis symbols of ions. |NaCl or MgS |

|Lattice Energy: Measurement of the strength of the ionic bond | |

|(Hlattice = energy required to completely separate 1 mole of solid ionic compound into its gaseous | |

|ions |Covalent bond strength is measured by bond energy. Bond energy is|

|Electrostatic attraction (and thus lattice energy) increases as ionic charges increase and as ionic|calculated by |

|radii decrease. |= Energy of reactant bonds- Energy of product bonds |

|[pic] | |

|Covalent Bonding: Atoms share electrons to achieve noble gas configuration that is lower in energy | |

|(therefore more stable) | |

|Occurs when difference of electronegativity is ≤ 1.7 | |

|Polar covalent: 0.3 < diff in EN ≤ 1.7 | |

|Nonpolar covalent: 0 ≤ diff in EN ≤ 0.3 | |

|Coordinate covalent: shared pair contributed by only one of the two sharing species. Ex: Lewis | |

|acids and bases | |

|Metallic Bonding: Forms between metal atoms because of the movement of valence electrons |[pic][pic] |

|from atom to atom to atom in a “sea of electrons”. The metal consists of cations held |[pic] |

|together by negatively charged electron “glue”. | |

|Results in excellent thermal and electrical conductivity, ductility, and malleability | |

|A combination of 2 metals is called an alloy | |

|The Octet Rule: Atoms tend to gain, lose, or share electrons until they are surrounded by 8 | |

|electrons in their outermost energy level (filled s&p shells), and are thus energetically | |

|stable | |

|Lewis symbols (electron-dot symbol) | |

|Shows a dot for valence electrons of an atom/ion | |

|Places dots at top, bottom, right, and left sides and in pairs only when necessary (Hund’s | |

|rule). | |

|Primarily used for representative elements only (Groups 1A – 8A) | |

|Transition metals typically form +1, +2, & +3 ions | |

|Transition metal atoms first lose both “s” electrons, even though it is a higher energy | |

|subshell. Ex. Cr2+, Cr3+ |Ex. Draw the Lewis structures for the following using steps. Show work! |

|Most lose electrons to end up with a filled or half-filled subshell. Ex. Cu+ ion |Cl2 |

|Lewis structures: used to depict bonding pairs and lone pairs of electrons in the molecule |CH2Cl2 |

|Total the # of valence electrons in the system. Add the total negative charge if you have an|NH3 |

|ion, subtract the charge if you have a cation. |NaCl |

|Number the electrons if each atom is to be “happy” (8 electrons for octet rule, or 2 for |HCN(SO4) |

|hydrogen) |H2O2CNS |

|Calculate the number of bonds in the system. Covalent bonds are made by the sharing of | |

|electrons. # of bonds = (electrons in step 2 – electrons in step 1) / 2 | |

|Draw the structure. The central atom is usually the atom with the least electronegativity. | |

|Double check your answer by counting total number of electrons | |

|Exceptions to the Octet Rule | |

|Odd-electron molecules. Ex. NO or NO2 | |

|Incomplete octet. Ex. H2, He, BeF2, BF3 | |

|Expanded octet (occurs in molecules when the central atom is beyond the third period. The | |

|empty 3d subshell is used in hybridization). Ex. PCl5, SF6 | |

|Formal Charge: the numerical difference between # of valence electrons in the isolated atom | |

|and # of electrons assigned to that atom in the Lewis structure. Doesn’t represent real |# Valence electrons in free atom |

|charges, just a useful tool for selecting most stable Lewis structure |– # Non-bonding electrons |

|Assign unshared electrons (usually in pairs) to the atom on which they are found |– ½ (# Bonding electrons) |

|Assign one electron from each bonding pair to each atom in the bond (split the electrons in |Formal Charge |

|a bond) | |

|Subtract the electrons assigned from the original number of valence electrons | |

|Used to select the most stable (and therefore most likely) structure when more than one | |

|structure is reasonable according to rules |Ex. Draw at least 2 Lewis structures for each, then calculate the formal|

|Most stable: |charge for each atom in each structure. |

|Has FC on all atoms closest to zero |SCN- |

|Has all negative FC on most EN atoms |N2OBF3 |

|Resonance Structures: Equivalent Lewis structures that describe a molecule with more than | |

|one likely arrangement of electrons | |

|Notation: use double-headed arrow between all resonance structures |Ex. Draw the Lewis structure and determine the bond S-O, C-C, and C-H |

|“Real” electron structure of the molecule is an “average” of all resonance structures |bond orders. |

|Bond Order: Indication of bond strength and bond length |SO3 |

|Bond Enthalpy: Amount of energy required to break a particular bond between two elements in |C6H6 |

|gaseous state. Indicates “strength” of a bond. | |

|(Hrxn ≈ ( (Hbonds broken) - ( (Hbonds formed) | |

| |Ex: CH4 (g) + Cl2 (g) → CH3Cl (g) + HCl (g) DHrxn = ? |

| |Bond Ave DH/mol Bond Ave DH/mol |

| |C-H 413 Cl-Cl 242 |

| |H-Cl 431 C-Cl 328 |

| |C-C 348 C=C 614 |

| | |

| |Ans: (Hrxn ≈ -104 kJ/mol |

|Chapter 9 Summary Notes |

|Main Concepts |Explanations |

|Valence Shell Electron Pair Repulsion Theory: helps construct molecular 3-D shape | |

|from 2-D Lewis Structures |[pic] |

|Electron Domains: areas of valence electron density around the central atom | |

|Includes bonding electron pairs and lone electron pairs |[pic] |

|A single, double, or triple bond counts as one domain | |

|Basis for VSEPR: each group of valence electrons (electron domains) around a |[pic] |

|central atom tend to be as far as possible from each other to minimize repulsion | |

|and this determines molecular geometry of molecule | |

|lonepair-lonepair repulsion > lonepair-bondpair repulsion > bondpair-bondpair | |

|repulsion | |

|Molecular Dipole Moment: molecules with polar covalent bonds might have a net |Draw hybrid orbital diagram for C2H4. |

|dipole moment depending on the compound's 3-D geometry and symmetry |[pic] |

|Check if individual bonds are polar and if individual dipole moments cancel out | |

|due to symmetry | |

|Polarity: polar substances are soluble in water and non-polar substances are | |

|soluble in non-polar solvents like benzene and oil | |

|Valence Bond Theory: predicts bond strengths based upon orbital overlap for | |

|covalent bond formation | |

|Basis for VB: covalent bond forms when orbitals of two atoms overlap | |

|An orbital can have a max of two electrons with opposing signs | |

|The bond strength depends on the attraction of nuclei for the shared electron, so | |

|greater the overlap, the stronger the bond | |

|Sigma Bond: end to end overlap of orbital which allows free rotation of parts of | |

|molecule (single bonds) | |

|Pi Bond: side to side overlap of orbital which restricts rotation (2nd and 3rd | |

|bonds in double and triple bonds) | |

|Hybrid Orbital Theory: an extension of VB theory, where atomic orbitals "hybridize" to form |[pic] |

|new hybrid orbitals; explains the bonding in terms of quantum mechanical model of atom (s, |[pic] |

|p, d, f orbitals) |[pic] |

|Basis for HO: valence atomic orbitals in the molecule are very different from those in | |

|isolated atoms |[pic] |

|The number of hybrid orbitals obtained equals the number of atomic orbitals mixed | |

|The type of hybrid orbitals obtained varies with types of atomic orbitals mixed | |

|ns and np give two sp hybrids | |

|ns and two np give three sp2 hybrid orbitals | |

|ns and three np give four sp3 hybrid orbitals | |

|Delocalized Bonds are present in compounds with resonance structures | |

|Molecular Orbital Theory: is based on the wave nature of the electrons and is a better model| |

|to explain paramagnetism of oxygen | |

|Basis for MO: Molecular orbitals form through the combination of atomic orbitals | |

|Bonding MO: stable orbital that forms between nuclei | |

|Antibonding MO: less stable orbital that forms behind nuclei | |

|Sigma MO: orbital forming from a combination of two 1s or 2s orbitals form different atoms | |

|or two 2pz orbitals from different atoms | |

|Pi MO: orbital forming from a combination of two 2px or 2py from different atoms (do not | |

|appear until B2) | |

|Diamagnetism: all electrons paired; no magnetic properties | |

|Paramagnetic: at least 1 unpaired electron; drawn into exterior magnetic field since spins | |

|of atoms become aligned; unlikely to retain alignment when field is removed | |

|Ferromagnetism: occurs primarily in Fe, Co, and Ni; drawn into exterior magnetic field since| |

|spins of atoms become aligned; likely to retain alignment when field is removed | |

|Chapter 10 Summary Notes |

|Main Concepts |Explanations |

|Characteristics of Gases | |

|Particles in a gas are very far apart, and have almost no interaction. | |

|Gases expand spontaneously to fill their container (have indefinite | |

|volume and shape.) | |

|Pressure | |

|Pressure = a force that acts on a given area | |

|[pic] | |

|Atmospheric pressure: the result of the bombardment of air molecules | |

|upon all surfaces | |

|1 atm = 760 mm Hg = 760 torr = 101.3 kPa = 14.7 PSI | |

|Barometer: measures atmospheric P compared to a vacuum | |

|* Invented by Torricelli in 1643 | |

|* Liquid Hg is pushed up the closed glass tube by air pressure | |

|Manometers: measure P of a gas | |

|*Closed-end: difference in Hg levels (Dh) shows P of gas in container | |

|compared to a vacuum | |

|* Open-end: Difference in Hg levels (Dh) shows P of gas in container | |

|compared to Patm | |

|Gas Laws | |

|Boyle’s Law: the volume (V) of a fixed quantity (n) of a gas is | |

|inversely proportional to the pressure at constant temperature (T). | |

|Main Concepts |Explanations |

|Charles’ Law: V of a fixed quantity of a gas is directly | |

|proportional to its absolute T at constant P. | |

|Gay-Lussac’s law: P of a fixed quantity of a gas is directly | |

|proportional to its absolute T at constant V. | |

|Seen as derivative of C’s and B’s laws | |

|Avogadro’s hypothesis: Equal volumes of gases at the same T & | |

|P contain equal numbers of molecules | |

|Combined Gas Law | |

|Chapter 11 Summary Notes |

|Main Concepts |Explanations |

|INTRAmolecular Forces: the forces holding atoms together to form moleculs | |

|INTERmolecular Forces: Forces between molecules between ions, or between | |

|molecules and ions | |

|Intermolecular Forces (IMF) | |

|-IMF < intramolecular forces (covalent, metallic, ionic bonds) |Boiling points and melting points are good indicators of relative IMF strength |

|-IMF strength: solids>liquids>gases | |

|-Types of IMFs Ion-Ion Forces, Ion-Dipole Forces, Dipole-Dipole Forces, H-bonds | |

|extreme dipole-dipole, LDFs | |

|Types of IMF | |

|Electrostatic Forces: act over larger distances in accordance with Coulomb’s Law| |

|Ion-Dipole: between an ion and a dipole (a neutral, polar molecule has separate | |

|partial charges) | |

|-Increasing with increasing polarity of molecule and increasing ion charge | |

|Ion-Permanent Dipole | |

|-Water is highly polar and can interact with positive ions to give hydrated ions| |

|in water | |

|-Attraction between ions and dipole depends on ion charge and ion-dipole | |

|distance | |

|Dipole-Dipole | |

|-Weakest electrostatic force (not all IMFs, LDFs weaker than dipole-dipole); | |

|exist between neutral polar molecules | |

|-Increase with increasing polarity (dipole moment) of molecule | |

|Hydrogen Bonds (H-bonds) | |

|-H is unique among elements because it has a single e- that is also a valence e-| |

|-When e- is “hogged” by a highly electronegative atom (very polar covalent | |

|bond), the H nucleus is partially exposed and becomes attracted to e- rich atom | |

|nearby | |

|-Explains why ice floats on water, has lattice-like structure, explains why | |

|molecules with H-bonds have higher boiling points, H-bonding in water, O -H bond|[pic][pic] |

|is very polar |[pic] |

|H-bonding in biology | |

|-DNA bases bind to each other due to specific hydrogen bonding between Lewis | |

|Bases | |

| | |

| | |

| | |

| | |

|Inductive Forces arise from induced distortion of e- cloud | |

|London Dispersion: between polar or nonpolar molecules or atoms, but is | |

|generally mentioned for non polar molecules when other forces are absent. Very |[pic] [pic] |

|weak, motion of e- creates an instantaneous dipole moment which induces a dipole|Molecular |

|in an adjacent atom |[pic][pic][pic] |

|Nonpolar molecules can dissolve in water due to LDFs. Water induces a dipole in |Ionic Covalent Network Metallic |

|electron cloud. Solubility increases with mass of gas due to greater distortion.|Credits: Google Images |

|When induced forces between molecules are very weak, the solid will sublime | |

|(solid to gas) | |

|Liquids | |

|Molecules are in constant motion, molecules close together | |

|Liquids are almost incompressible | |

|Evaporation | |

|-To evaporate, molecules must have sufficient energy to break IMFs | |

|-Condensation is reverse (remove energy and make IM bonds) | |

|Vapor Pressure | |

|Heat of Vaporization heat required (at constant P) to vaporize the liquid | |

|Equilibrium vapor pressure & the Clausius-Clapeyron Equation | |

|Used to find ∆vapH˚ | |

|Logarithm of vapor pressure P is proportional to ∆vapH˚ and to 1/T | |

|lnP = -(∆vapH˚/RT) + C | |

|Surface Tension leads to spherical liquid droplets | |

|Properties resulting from IMFs | |

|Viscosity: resistance of a liquid to flow | |

|Surface Tension: energy required to increase the surface area of a liquid | |

|Intermolecular forces lead to capillary action and concave meniscus for a water | |

|column | |

|-Capillary Action: movement of water up a piece of paper depends on the H-bonds | |

|between H2O and the OH groups of the cellulose in the paper | |

| | |

|Cohesion: attraction of molecules for other molecules of the same compound | |

|Adhesion: attraction of molecules for a surface | |

|Meniscus: curved upper surface of a liquid in a container; a relative measure of| |

|adhesive and cohesive forces | |

|London Dispersion Forces | |

|-Increase with increasing molar weight, increasing # of e-, increasing # of | |

|atoms | |

|-“Longer” shapes (more likely to interact with other molecules) | |

|Phase Changes | |

|-Endothermic: melting, vaporization, sublimation | |

|-Exothermic: condensation, freezing, deposition | |

|Structures of solids | |

|Amorphous: without orderly structure (ex. Rubber, glass) | |

|Crystalline: repeating structure; have many different stacking patterns based on| |

|chemical formula, atomic or ionic sizes, and bonding | |

|Types of Crystalline Solids: | |

|-Atomic: Properties: poor conductors, low melting point | |

|-Molecular: Properties: poor conductors, low to moderate melting point | |

|-Ionic: Properties:hard and brittle, high melting point, poor conductors, some | |

|solubility in H2O | |

|-Covalent (a.k.a. covalent network): Properties: very hard, very high melting | |

|point, generally insoluble, variable conductivity | |

|-Metallic: Properties: excellent conductors, malleable, ductile, high but wide | |

|range of melting points | |

|Chapter 13 Summary Notes |

|Main Concepts |Explanations |

|Vocabulary |Describe the steps with the lab equipment needed to make 0.1 M 100 ml NaOH? (Hint: volumetric flask) |

|Molarity: measure of concentration in solutions, mol/L | |

| |How are supersaturated solutions created? |

|Solvation: dissolving; the interactions between solute and |Dissolve solute with heat, then cool solution slowly. Solution is “tricked” into appearing |

|solvent |unsaturated |

|Crystallization: process by which solute particles leave | |

|solvent |At 40oC, the solubility of KNO3 in 100g of water is 64 g. |

| | |

|Saturation |What is the solubility of KCl at 10oC? |

|Saturation: solution that is in equilibrium with undissolved |30 g |

|solute | |

|Unsaturated solution contains less solute than saturated |At which points would an unsaturated solution appear? Supersaturated solution? |

|solution | |

|Supersaturated solution contains more solute than saturated |How much KClO3 needs to be added to 10g of KClO3 at 60oC to make a saturated solution? |

|solution but appears unsaturated |20 g |

| | |

|Factors Affecting Solubility | |

|Miscible liquids mix; both are polar or both are nonpolar | |

|Covalent network solids do not dissolve in polar or nonpolar |What is normal boiling point of ethanoic acid? |

|solvents |117 oC |

|Increasing temperature increases solubility for most solids, | |

|but decreases solubility for gases | |

| | |

|Chapter 14 Summary Notes |

|Main Concepts |Explanations |

|Reaction rate: A measure of the (average) speed of a reaction |Rate with regards to other |

|Reaction rate is affected by: |reactants and products |

|Concentration of reactants |2N2O5 (g) → 4 NO2 (g) |

|Temperature of the reaction |+ O2(g) |

|Presence/absence of a catalyst |If D[O2]/Dt = 5.0 M/s, what is D[N2O5]/Dt?-10.0 M/s |

|Surface area of solid or liquid reactants and/or catalysts | |

|-Average Rate: Rate of a reaction over a given period of time |Orders of Reactions & Related Equations: |

|-Instantaneous Rate: Rate of reaction at ONE given point of |Using Initial Rates Method: data given |

|time. |Experiment |

|-Initial Rate: Rate of reaction at t=0. (its instantaneous rate|[A] |

|at t=0) |[B] |

|Reaction order: the exponents in a rate law (can be fractions) |Initial Rate of Formation of |

|Rate law: shows how the rate of reaction depends on the |C in M |

|concentration of reactant(s); determined experimentally. Cannot| |

|be determined by the coefficients of a balanced reaction |1 |

|(unless in an elementary step) |0.60 |

|Rate = k[A]m[B]n |0.15 |

|*Units of k change w/order of the rxn |6.3´10-3 |

|To Find Rate Laws: | |

|Using Initial Rates |2 |

|Integration Method: Determining Rate Law by Determining the |0.20 |

|Change in Concentration of reactants over time |0.60 |

|gives rate law either graphically or by calculations. |2.8´10-3 |

| | |

| |3 |

| |0.20 |

| |0.15 |

| |7.0´10-4 |

| | |

| | |

| |Integration Method (Using graphs or equations) |

| |Zero Order |

| | |

| | |

| | |

| | |

| | |

| |[A]t = -kt + [A]0 |

| |(k) = M/s |

| | |

|The Collision Model | |

|-Reactants must collide, and with the right orientation and | |

|energy for an effective collision | |

|-Elementary steps: a single event or step (reaction) in a |First Order |

|multi-step reaction |[pic] |

|-Molecularity: the # of molecules participating as reactants in|ln[A]t = -kt + ln[A]0 or log[A]t = -kt / 2.303 + log[A]0 |

|an elementary step |(k) = s-1 |

|-Catalyst: Substance that changes the rate of a reaction | |

|without undergoing a permanent chemical change itself | |

| | |

|Check for Permissible Rxn. Mechanism | |

|Balanced eq. | |

|Rate Determining Step (RDS) is the slow one. | |

| | |

| | |

|-Radioactive decay follows a first order kinetics. Half life | |

|(t1/2) can be calculated, if rate constant is known or vice | |

|versa. | |

| | |

| | |

|-Activation energy ( Ea ): minimum energy required to initiate |Second Order |

|a chemical reaction | |

| |[pic] |

| | |

| | |

| | |

| | |

| | |

| | |

| | |

| | |

| | |

| |Half Life |

| | |

| | |

| | |

| | |

| | |

| |Activation Energy |

| |[pic] |

|Chapter 15 Summary Notes |

|Main Concepts |Explanations |

|Chemical Equilibrium | |

|Occurs when rate of forward reaction = rate of reverse reaction. Ex. Vapor pressure: | |

|rate of vaporization = rate of condensation, Saturated solution: rate of dissociation| |

|= rate of crystallization. Conc. of reactants and products does not have to be equal | |

|at the equilibrium, only the rates of forward and reverse rxn become equal. | |

|Express concentration in Partial Pressure for gases and molarity for solutes in | |

|liquids | |

|Rate = kforward [A] | |

|Rate = kreverse [B] | |

|[pic] at equilibrium | |

|-If Kc > 1, then more products at equilibrium | |

|-If Kc < 1, then more reactants at equilibrium | |

|-If Kc = 1, then almost equal concentrations of products and reactants | |

|There is a spontaneous tendency towards equilibrium. |The equilibrium expression is: |

|(spontaneous ≠ quickly, spontaneous = always moving towards equilibrium) | |

|It is possible to force equilibrium one way or the other temporarily by altering the | |

|reaction conditions, but once this “stress” is removed, the system will return to its|For a heterogeneous equilibrium: |

|original equilibrium. |CaCO3 (s) ↔ CaO (s) + CO2 (g) |

|Law of Mass Action : | |

|a A + b B ↔ c C + d D | |

| | |

|Concentrations of pure solids and pure liquids are not included in Keq | |

| | |

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Summary of the Chapter and Important things to remember:

Summary of the Chapter and Important things to remember:

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Summary of the page and Important things to remember:

Summary of the page and Important things to remember:

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Summary of the page and Important things to remember:

Summary of the page and Important things to remember:

Zeff = Z − S

Zeff = ENC

Z = atomic number

S = screening constant (# of inner electrons)

Summary of the Chapter and Important things to remember:

Summary of the page and Important things to remember:

Summary of the page and Important things to remember:

Summary of the page and Important things to remember:

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Summary of the Chapter and Important things to remember:

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Summary of the page:

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Summary of the Chapter and Important things to remember:

Summary of the Chapter and Important things to remember:

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1 = kt + 1

[A] [A]0

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For a graph between ln k and 1/T, the slope can be used to calculate Ea—activation energy. Remember to use 8.314 j/mol. K for R.

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Summary of the Chapter and Important things to remember:

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