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:
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:
Summary of the page:
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:
Summary of the page and Important things to remember:
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