Nikhil Kanthi Period 2 Semester 1
Chemistry Honors Reading Notes Compilation
Chapter 1 Notes 1. Chemistry is study of substances and their change 2. Scientific Method a. Make observation/ask question b. Make hypothesis c. Test hypothesis d. Answer will explain natural law (how), form natural theory on why 3. Safety a. If it is hot, let it cool b. Avoid awkward transfers 4. Measurement a. SI system (exclude Celsius and Liters) b. Final digit is always estimated c. Precision vs Accuracy i. Precision is same answer again and again ii. Accuracy is right answer again and again 5. Sig Figs a. Pacific Atlantic Rule (Pacific for Decimals, Atlantic for no decimals) i. Multiplication-division have least amount of sig figs in answer ii. Addition-Subtraction have highest uncertainty b. Percent Error --> measured - accepted/accepted 6. Dimensional Analysis a. Converting units by use of fractions
Chapter 2 Notes 1. Energy b. Capacity to do work i. Radiant energy (sun) ii. Kinetic energy (motion) iii. Potential Energy (gravity) --> stored c. Measured in calories, Calories is 1000 calories (used in food) (1 cal -->4.184 J) d. Law of Conservation of Energy i. Different forms of energy can be converted to each other without wasting any energy in between e. Different Units of Energy i. Fahrenheit --> first one (-32, x5, /9) ii. Celsius --> 0 is freezing water, 100 at boiling iii. Kelvin --> same unit as Celsius, 0 in Kelvin is 273 in Celsius 2 . Matter f. Liquid changes to suit volume, gas expands to suit volume, solid does nothing g. Physical Properties i. Can be observed h. Chemical Properties i. Can be observed by changing matter i. Law of Conservation of Mass i. Cannot add or subtract mass j. Elements i. Substance that cannot be split into new substances ii. Two or more elements make a compound 1. Compound name is element names with subscripts for amount k. Mixtures i. Blend of two or more substances ii. Visibly distinguishable --> heterogenous 1. Separated by filters l. Homogenous --> not visibly distinguishable i. Separated by 1. Distillation a. Differences in boiling points means one evaporates, leaving the other 2. Crystalization a. Partial evaporation so crystals will form 3. Chromatography a. Letting it flow on solitary substance that collects it
Chapter 5-6 Reading Notes 1. Chemicals organized via periodic table a. Dobereiner --> three groups by properties i. Middle group’s atomic mass --> average of other two masses b. J.A.R. Newlands --> grouped by 8, repeating patterns c. Mendeleev --> Periodic cards by atomic mass, patterns emerged i. Said some were calculated incorrectly d. H.G.J. Moseley --> metals hit with electrons make x-rays of different frequencies because different positive charge (atomic number) i. That’s how to arrange it by e. Periodic Law --> elements arranged by atomic number, patterns become obvious 2. Parts to Periodic Table a. American system is by tall columns (1A, 2A) and short columns (1B) b. Families hold traits, periods rise by electrons c. Period is energy level, Column is number of valence i. Metals--> malleable, ductile, good conductors, mostly solid at room temperature ii. Nonmetals --> nonmalleable, not ductile, bad conductors, mostly gaseous at room temperature, iii. Metalloids --> little of both d. Electron configuration gives you patterns i. Vertical gives you same energy level, similar traits ii. Valence electrons responsible for traits, same valence number in family 3. Blocks a. S-blocks are alkali, alkali earth b. P-blocks are gaseous with oxygen, halogens, noble gases c. D-blocks are 10 elements long with all metals d. F-blocks are 14 elements long with all metals e. S and P are representative, D and F are transition metals 4. Periodic Trends a. Atomic radius increases down (new energy level) decreases sideways (positive protons want electrons near them) b. When atom loses electron, protons pull other electrons closer (smaller), when atom gains one electrons repulse it (bigger) c. Elements of a same family have same-charge ions, right side is negative, left is positive d. Ionization Energy --> energy to remove one electron 1. measured by moles (6.02 x 10^23) ii. Ionization decreases as you go down (more willing to lose, repulsion), increases as you go sideways (positive protons pull electrons to them) e. First Ionization Energy --> energy to remove one electron i. Jumps up after has full valence electrons f. Electron Affinity --> energy change in atom when atom gains electron
i. Most have a negative, release energy to gain electron g. Oclet Rule --> element’s tend to gain or lose electron to have full set of 8 h. Electronegativity --> want to gain electron (decrease down, increase right) 5. Groups and Tendencies a. Alkali Metals (not found alone in nature) i. malleable, ductile, good conductors (cut with knife) ii. intensely reactive (with halogens) because have 1 valence electron b. Alkali Earth Metals i. higher density and melting points than alkali metals ii. not as reactive, still not found alone in nature a. 1 metal + other metal --> alloy 2. Reacts with water/steam c. Transition Metals i. Chromium 1. Resistant to Corrosion 2. Used in many alloys (stainless steel) ii. Iron 1. Least expensive of all transition metals 2. Used in making alloys (steel) 3. 4th most abundant metal (obtained via blast furnaces) from Fe2O3 d. Coinage Metals i. Copper --> soft red metal, coinage, electrical wiring, (tin+copper-->bronze) ii. Silver --> white ductile malleable metal, best electrical conductor, too soft, used in alloy iii. Gold --> dense, soft, used in karats (fraction of karat/24 --> percent of gold in it) e. Inner Transition Metals i. 4f called Lanthanides, 5f called Actinides ii. Do not follow similar trends 1. Lanthanides all want to lose 3 electrons, radioactive, occur together in nature 2. Actinides --> radioactive f. Boron group --> aluminum i. Third most abundant on earth, found in compounds w/ silicon and oxygen ii. Low density but forms strong alloys g. Carbon Group i. Carbon --> (nonmetal) found in limestone, various densities (plant, diamonds) 1. Hydrocarbons --> carbons + hydrogens a. When burned with less or more oxygen, make carbon monoxide and dioxide respectively ii. Silicon 1. 2nd most abundant (used in glass, found in sand) 2. Pure silicon used in chips h. Nitrogen Group
i.
j. k.
l.
i. Nitrogen --> 80% of atmosphere (found as N2) 1. Essential for life, not very reactive, nitrogen fixation makes it into compounds a. Ammonia ii. Phosphorus --> compounds called phosphates 1. Helps in fertilizer Oxygen Group i. Oxygen --> Most abundant on planet (89% water, 23% air, 46% rocks) 1. Found in atmosphere as O2 and O3 ii. Sulfur --> found as fools gold, characterized by odor, used in sulfuric acid Halogens --> salt formers, highly reactive (fluoride is corrosive gas) i. Chlorine is used to kill bacteria, found in bleach Noble Gases i. Argon found first, new group must exist, very rare on the planet 1. Reactivity Increases with Size Hydrogen --> Most abundant in universe, found as H2, light enough to escape gravity, used as ammonia
Chapter 7-8 Notes 1. Ionic Bonds a. Positively charged atom stuck to a negatively charged atom b. Ionic compound i. High melting points ii. Brittle iii. Water soluble iv. Liquid form conducts electricity c. Positive charge (cation) and negatively charged (anion) 2. Octet rule a. to move to nearest noble gas position (left for anion, right for cation) i. Transition metals do not follow pattern 3. Lewis Dot Diagram a. 2 Dots on each side around element symbol symbolizes valence electron count 4. Monatomic and Polyatomic a. Monatomic cations (ions formed from one element [Na+, O2-) i. Alkali family ions (1+) ii. Alkali earth family (2+) b. Metals make cations of more than one type i. Fe2+, Fe3+ 1. Differentiation via Roman numerals c. Monatomic Anions i. Mostly nonmetals 1. Labeled with suffix ‘ide’ d. Polyatomic Ions i. Made internally with covalent bonds (sulfate) 1. Bond with with ions by ionic bonds 5. Binary Ions and Notations a. Binary Ionic Compounds --> Cations then anions b. Empirical Formula --> simplest ratio of atoms to atoms i. H2O (not H4O2) c. Crisscross Method i. Take individual ion charges and put them as subscript notations ii. For polyatomic ions, put subscript after parenthesis around polyatomic ion 6. Covalent Bonds --> shared electrons, not given, to get octet satisfied a. atoms grouped by covalent bonds are called molecules b. Molecular formula (like empirical, but not clean ratios) i. Different combinations of same ratio of elements gives different elements 1. C6H12O6 --> glucose, C3H6O3 --> lactic acid c. Structural formula is used i. Many different kinds, common is Lewis Structure
7. Lewis Structure formula a. Like Lewis Dot, but all elements together in molecule represented i. Shared electrons circled 8. Single, Double, Triple Bonds a. Single --> One shared pair between two elements b. Double --> 2 pairs shared c. Triple --> 3 Pairs Shared d. No Quadruple, because full 8 electrons would be shared 9. Dash Formula --> Dash instead of two dots for shared electrons 10. There are some exceptions to octet rule (boron and others make more or less that octet electron sets) 11. Polarity a. Electronegativity (atoms attraction to shared electrons) i. When one specific part of molecule is more negative, molecule is polar b. Changes in charge of electronegative electron is called partial change c. If equal electronegativity, then nonpolar i. If electronegative atom at center, it is also nonpolar d. Difference in electronegativities i. Less than 0.4 --> nonpolar covalent ii. 0.4 --> 2 --> polar covalent iii. Greater than 2 --> ionic 12. Naming Compounds a. Named by bonds and atoms b. Ionic Compounds i. cation first [as in chemical formula]) ii. To balance the compound (for metal) use Roman Numerals c. Hydrates i. Absorb water into their solid structures 1. Anhydrous substances are hydrates without their water ii. Name is name of anhydrous substance and then prefix-hydrate 1. pentahydrate, octahydrate, trihydrate iii. In chemical notation, it is times sign and then hydrate with number before it d. Molecular Compounds i. Prefix for amount of atom in molecule (like hydrate notation) ii. Suffix -ide added to more electronegative element 1. Prefix mono not for first one, oo’s not used, common names used if there e. Acids i. Hydrogen + Anion 1. If Anion ends in -ide, then it is hydro+anion ending with ic and then acid ii. If Anion does not end in -ide, then it is anion name with ic and then acid 13. Structural Formula shows bonds, not bond shape a. Ball and Stick model
b. 14. a. 15. a. b. c. 16. a.
b.
c.
d.
e.
17. a.
18. a. b. 19.
i. Ball is nucleus, sticks are one bond or two bonds (straight or curvy) ii. Used to give 3D representation of atom All shapes are symmetrical, electrons repulse each other VSEPR Electrons repulse each other because of charge i. Electrons placed as far from each other as possible Molecular Shape and Categories Bond Angle Shape Bond Polarity Molecular Shapes Linear i. Molecules arranged in a straight line 1. All diatomic molecules ii. 180 degrees bond angle Trigonal Planar i. Flat and triangular shape ii. Two bonds are single, one is double iii. 120 degrees Tetrahedral i. 109.5 degrees ii. One central atom and four around it (normally) and no lone pairs Trigonal Pyramidal i. Three pairs and one lone pair ii. 107 degrees iii. Lone pair takes up more space because has higher repulsion Bent i. 105 degrees ii. Looks like linear, but 2 unshared pairs push it further, making it bent to a side Hybrid Orbitals When atoms come together, their orbitals morph i. Linear/Triple Bonds --> sp 1. When atom has two sets of atoms around it 2. Triple has it because at most triple bonds can exist, so only 2 atoms bonding ii. Trigonal Planar --> sp2 1. When atom has 3 sets around it (triangle) iii. Tetrahedral/Pyramidal/Bent --> sp3 1. When 4 sets around it (tetrahedron) Bond Length As you move down, bonds are longer, move across, bonds are shorter Multiple are shorter than singular Polarity
a. Shape and bonds determine polarity of a molecule b. All nonpolar charges means nonpolar c. Formaldehyde i. H2OC (all electrons go to O, so one end is charged, other isn’t) d. Carbon Dioxide i. Carbon Oxygen bond is polar ii. Bonds face in opposite direction, polarity is cancelled out iii. Since no polar and nonpolar end, cannot easily attract to another, and is a gas at room temperature e. Water i. Bent, so polar (all bent are polar) ii. Water is a liquid (part of almost every liquid on earth) because it can attract to other molecules 1. Hydrogen’s positive center attracts another oxygen’s negative, and chain is made f. Larger Molecules i. Polarity determines shape 1. Similar polarities attract one another 2. Shape can be determined by polarity
Chapter 12-13 1. All reactions need energy a. Breaking bonds requires, making releases 2. Thermochemistry is study of heat in chemistry (part of thermodynamics) 3. Exothermic vs Endothermic a. Products - Reactants b. Exothermic --> more heat released than absorbed c. Endothermic --> more energy absorbed than released 4. Total Energy --> kinetic and potential energy of particles a. Enthalpy is a little more i. Difference in enthalpy of products and reactants is energy b. Positive for endothermic, negative for exothermic c. Standard Enthalpy is (delta H with dot) 25 C and 1 atm (all substances need to be in pure form) i. If heat made in standard enthalpy goes over or below 25, enthalpy takes that into account d. Heat and moles are proportional 5. Hess’ Law a. Net enthalpy change = sum of individual steps needed in reaction b. If coefficients multiplied by factors, enthalpy multiplied too (heat and moles are proportional) c. If arrow is changed, sign is changed (endothermic becomes exothermic) 6. Calorimetry a. Heat flow and its measurement i. Delta H by delta H of surroundings (qrxn vs qsurr) b. Heat capacity --> energy to raise object by 1degree c. Specific heat --> raise by 1g of object by 1 degree d. Calorimeter --> insulated cup with known amount of water and thermometer i. If water heats up, exothermic, vice versa ii. Measured in J or kJ 7. Caloric Theory --> heat moved from hot things to cold things 8. Kinetic Theory --> heat is transfer of kinetic energy from hot to cold 9. Gases have uniform traits a. 22.4L at STP b. Mostly made of molecules (except for noble gases) c. Have mass and can be compressed d. Fill containers completely and exert pressure (this is dependent on temperature) e. Diffuse through others easily 10. Manometers --> tubes with open end and closed end to a closed container a. If gas on open end is lower, then atm is higher and inner pressure is less than atm, if it is higher, then inner pressure is higher and inner pressure is more than atm
11.
Kinetic Molecular Model a. Have mass, in constant motion, lot of space between them, collisions are perfectly elastic, heat makes them move faster 12. Gases described by amount (moles), volume (L), temperature (measured in C, displayed in K), and pressure (atm, kPa) 13. Gas Laws a. Boyle’s Law : Pressure and Volume are inverse (less pressure is more volume) b. Charles’ Law: Volume Changes according to Temp c. Avogadro’s Law: All gases at same temp and pressure occupy same volume d. Dalton: Sum of individual partial pressures is same as total pressures 14. Ideal Gas Law a. PV= nRT i. P is pressure, V is volume, n is moles, R = 0.0821 atm-L/molK, T is Kelvin/Celcius ii. R is constant (0.0821 atm, 8.134 at J and Pa) b. Exceptions at high pressure (individual volumes come into play) and at low temperatures (attractive forces come into play) 15. Density of gas depends on Temp, pressure, volume and molar mass, can be tweaked for less dense gas a. Helium safer than hydrogen, easier than hot air 16. Law Math Conversions and Mnemonic Devices a. P T V b. 1 atm = 101.3 kPa = 760.0 mmHg c. For R, convert all units to correspond with R (atm L / moles K) (Pa m^3/mol-K) (J/mol K)
Chapter 14-15 1. Liquids vs Solids vs Gases a. Gases take up whole container and match shape, liquids only match shape, solids retain shape size b. Liquids barely expand upon heating and barely compress (unlike gases) and move and diffuse slower than gases c. Solids have highest attraction forces and barely move d. All three’s state at room temp depends on intermolecular forces (weaker than other bonds) 2. Water at 0 C has too little kinetic energy to overcome attractive forces, above 0 C flow easily, above 100 C completely free of each other 3. Bond types a. Ionic --> holds two different elements together (solids at room temp), strong, pulled apart by water when they dissolve b. Metallic --> holds two metals together (solids at room temp), strong c. Covalent bonds --> molecules made of same elements (nonmetals) have intramolecular (between molecules) and intermolecular forces i. Boiling temp is a study of intermolecular forces (higher boiling means stronger forces) ii. Forces rise as you go down a family d. Intermolecular forces in molecules include dispersion (induced dipole makes other dipoles, very weak), dipole dipole (stronger, between permanent dipoles) and hydrogen bonds (strong dipole dipole due to low electronegativity of hydrogen atoms) 4. Liquids a. Viscosity is intermolecular force dependent function i. Resistance to friction, stronger intermolecular forces give higher viscosity ii. Decreases with temperature b. Surface tension --> unequally distributed forces of molecules on surface of liquid make surface taut, higher intermolecular forces increase surface tension 5. Solids a. Crystalline -->ordered unit cells (repeating pieces in larger structure) if solid has water, it’s hydrate b. Amorphous --> sometimes supercooled liquids, have no form, get softer as heated until liquid (ex: glass and rubber), think of them as liquids very cold with high viscosities c. Bonds in solids depend on intermolecular forces (hardness, melting points, electrical conducting) i. Metallic solids --> all ions together in sea of electrons, good conductors, malleable and ductile ii. Molecular solids --> weaker, cannot conduct, higher melting point --> higher intermolecular iii. Ionic --> strong forces of attraction, brittle, bad conductors (unless liquids), most stable iv.Covalent network --> very strong, made with covalent bonds, aren’t molecules
6. Changes of state and energy involved a. (Gas --> liquid --> solid) --> exothermic, <--- endothermic b. Vaporization --> making liquid into gas, is endothermic i. Anything over freezing is vaporizing, certain water particles are leaving surface to become gas ii. Volatile --> vaporizes fast iii. As they evaporate, solution cools because of less kinetic energy c. Condensation --> making gas into liquid (exothermic) i. Any particle hitting surface of liquid becomes liquid by intermolecular forces ii. In closed atmosphere, as more evaporate, more condense, until equilibrium is reached d. Vapor pressure --> pressure of vaporized molecules, at equilibrium pressure is constant i. Any substanceâ&#x20AC;&#x2122;s boiling point is at what temp vapor pressure is atmospheric pressure e. Freezing point --> melting point, not affected by external pressure, energy to melt is heat of fusion f. Sublimation and Deposition i. Sublimation --> solid to gas, opposite is deposition ii. Heat of sublimation = heat of deposition (opposite signs) iii. Close to heat of fusion+heat of vaporization g. Diagrams (phase diagrams and heating curves) i. Heating curves --> rises as heat is absorbed, constant at phase changes ii. For water, heat rises, and at 100 C, heat is kept on added until water completely evaporates, then temperature of steam rises until a certain point iii. Phase diagrams --> specific temperatures and pressures where two phases can occur iv.Triple point of substance is where all three states can be happening 7. Solution is homogenous mixture of two or more substances (particles are tiny, equally distributed, and solution will not settle if held in same conditions) a. Solute is broken down in solvent, solute is one who normally goes under change of phase b. Solid solutions are alloys, gaseous is air, liquid is just liquid (aqueous means solvent is water) i. If aqueous solutions have ions that can carry electricity, electrolytes, otherwise nonelectrolytes 8. Measuring Solutions --> concentration (how much solute in solvent) a. Molarity --> how much solute in solution (moles solute / L solution) b. Molality --> how much solute vs solvent (moles solute/kg solvent) c. Mole faction --> (moles of solute OR solvent / moles solution) how much of substance in solution d. Saturation --> it is saturated if maximum solute is in it, unsaturated if under maximum is in it, supersaturated if over maximum is in it (very unstable, adding any solute will make all solute crystals) 9. Solutions and Forces a. Intermolecular forces --> between solute and solvent
b. Solvation --> interaction between solutes and solvents (hydration if solvent is water) i. Ex: water molecules pull Na and Cl apart in NaCl in water ii. In solution, solutes and solvents are broken and intermingled (breaking bonds is endothermic, making bonds is exothermic) 10. Solubility --> how much solute will dissolve in solvent in specific conditions a. For solid solute, conditions are temperature, for gaseous, temperature and pressure b. Usually expressed per 100 g of solvent i. Solutes with polar molecules will dissolve in polar solvents (like dissolves like) ii. Polar polar, nonpolar-nonpolar, ionic-polar c. Temperature decreases gas’s solubility (more kinetic energy means escapes), increases solid’s solubility (most of the time) because more energy for solvent to dissolve them i. If temperature drops after adding solute, heating will help, if stays same heating won’t help, if increases, heating will decrease it d. Pressure --> helps gases significantly (more contact with surface of liquid) i. Henry’s Law --> solubility of gas is dependent on partial pressure of gases above it e. To dissolve faster (has nothing to do with solubility) i. Surface area of solute increase (more contact with solvent to pull apart) ii. Temperature (more kinetic energy for solvent) iii. Stirring (more undissolved solute contact with solvent) 11. Colligative Properties --> properties of solution based on concentration of solute, not identity a. Vapor Pressure Reduction --> nonvolatile solute makes less surface contact for solvent molecules to surface, less of them become gas (vapor pressure is proportional to concentration <-- Raoult’s Law) b. Boiling Point Elevation --> adding nonvolatile solute, less vapor --> more temp to make vapor pressure equal to atmospheric pressure (boil) i. Delta Tb is difference between boiling points, proportional to molality, constant and ifactor c. Freezing Point Depression --> lowering freezing point, same nonvolatile rules i. Also proportional to molarity and given constant d. Osmotic Pressure --> (solvent passes, solute does not) i. After certain difference on both sides of semi-permeable membrane, osmosis stops ii. Without osmotic pressure, solution is isotonic iii. With more solute, it is hypertonic (cell shrivel), with less solute, it is hypotonic (cell blows up) iv.Water follows solute e. Molar mass can be determined through any of the colligative properties
Chapter 16-17 1. Equilibrium --> two opposing processes occurring at the same rate a. Reversible reactions --> equilibrium where products become reactants while reactants become products i. Has to be reversible on its own ii. Forward reaction proceeds to the right, backwards proceeds to the left b. Reaction rate depends on concentration i. If products are taken away, reaction rate of forward reaction increases, vice versa c. Chemical Equilibrium --> When both rates are equal and the concentrations of reactants and products are constant 2. Equilibrium Constant a. Law of Mass Action --> Explains relative concentration of products and reactants in terms of equilibrium constant b. Equilibrium constant known as Keq --> concentration of products raised to coefficients over concentrations of reactants raised to coefficients i. Concentration --> molarity c. Law of Chemical Equilibrium --> every solution has a Keq (dependent on molarity) that varies by temperature i. You can read Keq --> if Keq >>1, then concentration of products is much greater than reactants at equilibrium, and if Keq <<1, then concentration of reactants is much greater than products d. You can measure the solution to check how close it is to equilibrium i. This is the equilibrium position of the solution e. Homogenous and Heterogenous Equilibria i. If it is homogenous, then does not change state, in heterogenous, changes state ii. If it changes to solid, omit concentration of solids in Keq f. Reaction Quotient --> uses specific data to see if reaction is at equilibrium i. If Q> Keq, then too many products, proceed to the left ii. If Q<Keq, then too many reactants, proceed to the right 3. Le Chatelierâ&#x20AC;&#x2122;s Principle a. Equilibrium systems can be altered at or nearing equilibrium b. If they are altered, the system changes to counter the change c. Concentration change --> if product or reactant is added, equilibrium will shift to other side to make more of other substance out of added substance d. Pressure change --> if more moles of product or reactant make less moles of opposite, then it will shift in that direction (more something used up to make less of something else) e. Temperature --> only one that changes Keq--> if reaction is exothermic or endothermic, then adding or subtracting heat will increase product or reactant, treat it like concentration of heat is being added f. Haber Process
i. Manipulating Le Chatelierâ&#x20AC;&#x2122;s Principle to make more ammonia by manipulating pressure and temperature 4. Dissolution and Precipitation a. In a solution, ionic solids separate into +/- ions i. Solution has to be a polar liquid b. Process of ionic solids becoming ions is called dissolution c. If one ion collides with the crystal, becomes part of the crystal i. If one ion touches another ion, both neutralize and become a solid d. Ions becoming solids again from ionic form is called precipitation i. Dissolution and precipitation are opposite systems e. After a while solubility is reached and no more ions can be dissolved i. Then, as more ions are dissoluting, some precipitate at the same time ii. When the rate of dissolution and precipitation is same, you have solubility equilibrium 5. Ksp --> Ksp is solubility product, is like Keq except no denominator a. No denominator because reactant is an ionic solid i. So it is product of productsâ&#x20AC;&#x2122; concentrations (at equilibrium) to powers of coefficients ii. Small value for Ksp means low solubility, vice versa b. When Ksp is given, can be used to find equilibrium concentrations (x^2, 4x^3, etc.) c. Ion Product --> input given values for concentrations and compare with Ksp (Q) i. If Q>Ksp, then ppt will form to get rid of excess ions, if Q<Ksp, no ppt ii. Ppt will keep on forming until Q=Ksp 6. Reactions that make precipitate are precipitate reactions a. All are double replacement reactions b. Precipitate must be a combination of ions (ppt must be ionic) c. Written as aqueous solutions on left, with ppt on the right (with the other aqueous product on right as well) i. Actual identity of precipitate can only be confirmed via experiment, no surefire way to predict d. All are double replacement because formation of ppt drives it further i. Solids have less energy than liquids, so tend to be more stable ii. Other driving forces are creation of water and gases 7. Ionic Equations a. Complete ionic equations have everything (all the spectator ions, those who do not undergo a change) i. When writing complete ionic equations, do not break up solids b. Net ionic equations exclude spectator ions 8. Common Ion Effect a. If two ionic substances sharing a common ion are mixed (one already in a saturated solution), increase of the added ion will result in ppt forming i. You are increasing the concentration of one ion past saturation (equilibrium), so ppt must form b. Also decreases solubility for newer ions who are introduced
Chapter 18-19 1. Properties of acids and bases a. Taste --> acids are sour and bases taste bitter b. Touch --> acids feel like water, bases feel slippery c. Reactions with metals --> acids react with mg, zinc, iron and aluminum d. Conductivity --> both are electrolytes (ionize in water), both good conductors e. Indicators --> change color in contact with them (ex: phenolphthalein, litmus paper) 2. Neutralization --> when acids and bases meet, cancel each other’s properties a. In proper ratio (MaVa=MbVb), no distinctive properties remain b. Products are ionic compound (salt) 3. Definitions of acids and bases a. Arrhenius --> (properties based on when acids/bases make contact with water) i. Acids release H+ ions (proton), bases release OH- ions ii. Acids + metals make H2 gas b. Bronsted Lowry definition i. Acid can donate H+ ion, base can accept H+ ion 1. Acids that can give 1 H+ are monoprotic, that can give 2 are diprotic, etc c. Conjugate acids and bases are when acids or bases lose/gain an H+ ion i. ex: HCl - H+ --> Cl- (<-- this is the conjugate base of HCl) d. Water is an amphoteric substance (can be an acid or a base in certain scenarios) i. Hydronium --> H+ ions in water actually make H3O+ ions 4. Strengths of acids and bases a. Strong acid readily makes H+ ions in water i. Indicated by ‘-->’ in chemical reactions, while weak ones have equilibrium arrows b. Strong bases readily accept H+ ions in water c. Strengths of acids/bases are inverse to their conjugates (strong acid=weak conjugate base) 5. Quantitative Formulas for Acids and Bases a. HA + H20 --> H30+ + Ab. Then Ka --> [H3O+][A-]/[HA] (refer to Mr. Cheung’s Ch 18 Notes for more in-depth info) i. For diprotic acids, more than 1 Ka c. Kb --> (same basic principle, replace HA with B and H3O+ with OH-) i. [HB][OH-]/[B] (refer to Mr. Cheung’s Ch 18 Notes for more in-depth info) d. Ka and Kb tell the strength of a base/acid (higher --> more disassociation in water --> more powerful) 6. Salt Hydrolysis --> salts made from neutralization reactions can be acidic or basic a. Strong acid + strong base --> neutral salt b. Strong acid + weak base --> acidic salt c. Weak acid + strong base --> basic salt d. Weak acid + weak base --> answers may vary 7. Naming acids
a. Acids = Binary acids --> hydrogen + element, Oxy acids --> hydrogen + oxygen + element, Carboxylic --> has carbon in it b. Bases = Bronsted Lowry always has one unshared pair of electrons i. Many anions are bases, amines are bases with 1 nitrogen atom with unshared pair c. Naming acids --> if name ends in i. ide, add ic to it and put after hydro, ii. in ate, add ic, no hydro, iii. in ite, add ous, no hydro 8. Self ionization of water --> even purest water has some OH- and H3O+ ions in it a. However, very little of of it is there i. At equilibrium (25 C), there are exactly 1.0 * 10^-7 M of both b. Kw --> 1.0 * 10^-14 i. This means that [OH-]x[H30+] = 1.0*10^-14 every time c. pH is -log[H3O+] i. More H3O+ means low pH, low pH means acidic, high pH means basic ii. Increase of pH by 1 is tenfold change d. pH found by indicators 9. Buffer --> can keep the pH of a solution constant even if something is added to solution a. Usually a mixture of acid/base and it’s conjugate make a buffer i. If acid is added, more H30+ ions will be made, then base/conjugate base in buffer will make H20 out of it (no change to pH) ii. If base is added, more HO- ions will be made, then acid/conjugate acid in buffer will make H20 out of it (no change to pH) b. Every buffer has a limit (when all of buffer’s acid/base and conjugate is used up) --> buffer capacity 10. Titration --> controlled neutralization reaction that lets you know concentration of acid/base a. Measure of pH is not the same as measure of concentration of acid b. For titration, need a standard solution (known concentration), unknown solution, and indicator c. After adding known amount of standard and a few drops of indicator, add unknown and at one point color of solution will change i. Point of color change is called end point (where more of unknown than standard in solution) d. Equivalence point --> when both are in same concentrations i. Equivalence point is very close to end point, so end point is appx equivalence point e. MaVa=MbVb, only unknown is Mb (moles of unknown), can be figured out f. There is a curve in pH as more unknown is added --> titration curve i. If curve is steep, then an indicator with a wide array of colors can work (sorry if this section is really vague, can’t write too concise and still keep stuff) g. For more info on titration and titration curves can be found on pgs 640, 641, and 642
Chapter 22-23 1. Chemical kinetics --> speed at which reactions occur a. Chemical reactions take time to occur b. Reaction rate is the change in concentration of reactants and products over a period of time i. Expressed in M/s c. Average rate of reaction --> (delta(reactants or products))/delta(time) i. Rate changes, so it makes sense to find average rate 2. Reaction mechanisms --> series of steps that lead from reactants to products a. Each step is an elementary step to the entire process b. Elementary steps must add up and (through canceling) equal the original reaction c. Some substances are made in elementary steps but cancel out i. These are intermediary products d. Speed of chemical reaction is dependent on slowest elementary step i. This step is called rate-determining step 3. Rate Laws --> equation that can be used to determine the rates of reactions a. Rate = k[A]x[B]y i. A and B are concentrations of reactants ii. x and y are powers of concentrations of reactants (must be found experimentally) 1. Not always the co-effecients iii. k is a fixed constant per reaction iv.Not all reactants are involved, those whose change in concentration do not affect reaction are excluded 4. Collision Theory --> two molecules have to collide to react with each other a. Only two can hit each other (very hard for three molecules to be at the same place) b. This is why elementary steps make sense i. If more than 2 reactants make more than 1 product, then various collisions need to take place for reaction to happen c. Not all collisions are effective, actually very few are effective i. Effective ones are called effective collisions, ineffective ones are ineffective collisions d. For collision to be effective, atoms must be in a position to break and form new bonds (orientation) and also both need to have enough energy to break and make bonds 5. Sufficient energy can be found from potential energy and kinetic energy a. Potential energy is energy waiting to be released, kinetic is released energy b. Energy required to break bonds comes from kinetic energy i. Kinetic energy is determined via mass and velocity of particle 6. Activation energy a. Energy of reactants must be raised to start reaction i. Difference between max energy of reactants and energy of products is activation energy b. Activation complex --> In transition state (brief moment where substance is neither product nor reactant) the substance is called activation complex
i. Has the most energy (at peak of activation energy) c. Activation energy (Ea) is energy required to make transition state and activation complex i. Activation complex can break up into products or reactants 7. Five factors affect reaction rates: nature of reactants, temperature, concentration, surface area, and catalysts a. Nature of reactants --> reactions that require slight rearrangement of reactants are faster i. Reactions that need many covalent bonds to be broken are slow ii. Reactions with reactants in different states are heterogenous (vice versa for homogenous) 1. Reaction between two gases will happen faster than two liquids or solids b. Temperature --> higher temperatures lead to faster reactions i. Molecules move faster with higher temperatures ii. Common rule of thumb --> increasing by 10 Celsius doubles or triples reaction rate c. Concentration --> more particles to collide means faster reaction d. Surface Area --> Larger surface area means more space for reactions to happen e. Catalysts --> Catalyst is a substance that increases the rate of reaction without being used up i. Catalysts in the body are called enzymes ii. Catalysts are opposites of intermediary products, are consumed first and then remade later iii. Catalysts lower activation energy, so more molecules can create products iv.Substances that reduce reaction rates are called inhibitors