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Nikhil Kanthi Period 2 Semester 1 Chemistry Honors Reading Notes Compilation

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An open source document I found that does a fairly good job of explaining the courseload of Chemistry Honors.


Page 1: Chemistry Honors Notes

Nikhil KanthiPeriod 2Semester 1

Chemistry Honors Reading Notes


Page 2: Chemistry Honors Notes

Chapter 1 Notes

1. Chemistry is study of substances and their change2. Scientific Method

a. Make observation/ask questionb. Make hypothesisc. Test hypothesisd. Answer will explain natural law (how), form natural theory on why

3. Safetya. If it is hot, let it coolb. Avoid awkward transfers

4. Measurementa. SI system (exclude Celsius and Liters)b. Final digit is always estimatedc. Precision vs Accuracy

i. Precision is same answer again and againii. Accuracy is right answer again and again

5. Sig Figsa. Pacific Atlantic Rule (Pacific for Decimals, Atlantic for no decimals)

i. Multiplication-division have least amount of sig figs in answerii. Addition-Subtraction have highest uncertainty

b. Percent Error --> measured - accepted/accepted6. Dimensional Analysis

a. Converting units by use of fractions

Page 3: Chemistry Honors Notes

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 Energyi. Fahrenheit --> first one (-32, x5, /9)ii. Celsius --> 0 is freezing water, 100 at boilingiii. Kelvin --> same unit as Celsius, 0 in Kelvin is 273 in Celsius

2 . Matterf. Liquid changes to suit volume, gas expands to suit volume, solid does nothingg. Physical Properties

i. Can be observedh. Chemical Properties

i. Can be observed by changing matteri. Law of Conservation of Mass

i. Cannot add or subtract massj. Elements

i. Substance that cannot be split into new substancesii. Two or more elements make a compound

1. Compound name is element names with subscripts for amountk. Mixtures

i. Blend of two or more substancesii. Visibly distinguishable --> heterogenous

1. Separated by filtersl. Homogenous --> not visibly distinguishable

i. Separated by1. Distillation

a. Differences in boiling points means one evaporates, leaving the other2. Crystalization

a. Partial evaporation so crystals will form3. Chromatography

a. Letting it flow on solitary substance that collects it

Page 4: Chemistry Honors Notes

Chapter 5-6 Reading Notes

1. Chemicals organized via periodic tablea. Dobereiner --> three groups by properties

i. Middle group’s atomic mass --> average of other two massesb. J.A.R. Newlands --> grouped by 8, repeating patternsc. Mendeleev --> Periodic cards by atomic mass, patterns emerged

i. Said some were calculated incorrectlyd. 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 obvious2. Parts to Periodic Table

a. American system is by tall columns (1A, 2A) and short columns (1B)b. Families hold traits, periods rise by electronsc. Period is energy level, Column is number of valence

i. Metals--> malleable, ductile, good conductors, mostly solid at room temperatureii. Nonmetals --> nonmalleable, not ductile, bad conductors, mostly gaseous at room tem-

perature, iii. Metalloids --> little of both

d. Electron configuration gives you patternsi. Vertical gives you same energy level, similar traitsii. Valence electrons responsible for traits, same valence number in family

3. Blocksa. S-blocks are alkali, alkali earthb. P-blocks are gaseous with oxygen, halogens, noble gasesc. D-blocks are 10 elements long with all metalsd. F-blocks are 14 elements long with all metalse. S and P are representative, D and F are transition metals

4. Periodic Trendsa. 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 positived. 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 electronsf. Electron Affinity --> energy change in atom when atom gains electron

Page 5: Chemistry Honors Notes

i. Most have a negative, release energy to gain electrong. Oclet Rule --> element’s tend to gain or lose electron to have full set of 8h. Electronegativity --> want to gain electron (decrease down, increase right)

5. Groups and Tendenciesa. 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 Metalsi. higher density and melting points than alkali metalsii. not as reactive, still not found alone in nature

a. 1 metal + other metal --> alloy2. Reacts with water/steam

c. Transition Metalsi. Chromium

1. Resistant to Corrosion2. Used in many alloys (stainless steel)

ii. Iron1. Least expensive of all transition metals2. Used in making alloys (steel)3. 4th most abundant metal (obtained via blast furnaces) from Fe2O3

d. Coinage Metalsi. Copper --> soft red metal, coinage, electrical wiring, (tin+copper-->bronze)ii. Silver --> white ductile malleable metal, best electrical conductor, too soft, used in al-

loyiii. Gold --> dense, soft, used in karats (fraction of karat/24 --> percent of gold in it)

e. Inner Transition Metalsi. 4f called Lanthanides, 5f called Actinidesii. Do not follow similar trends

1. Lanthanides all want to lose 3 electrons, radioactive, occur together in nature2. Actinides --> radioactive

f. Boron group --> aluminumi. Third most abundant on earth, found in compounds w/ silicon and oxygenii. Low density but forms strong alloys

g. Carbon Groupi. Carbon --> (nonmetal) found in limestone, various densities (plant, diamonds)

1. Hydrocarbons --> carbons + hydrogensa. When burned with less or more oxygen, make carbon monoxide and dioxide

respectivelyii. Silicon

1. 2nd most abundant (used in glass, found in sand)2. Pure silicon used in chips

h. Nitrogen Group

Page 6: Chemistry Honors Notes

i. Nitrogen --> 80% of atmosphere (found as N2)1. Essential for life, not very reactive, nitrogen fixation makes it into compounds

a. Ammoniaii. Phosphorus --> compounds called phosphates

1. Helps in fertilizeri. 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 acidj. Halogens --> salt formers, highly reactive (fluoride is corrosive gas)

i. Chlorine is used to kill bacteria, found in bleachk. Noble Gases

i. Argon found first, new group must exist, very rare on the planet1. Reactivity Increases with Size

l. Hydrogen --> Most abundant in universe, found as H2, light enough to escape gravity, used as ammonia

Page 7: Chemistry Honors Notes

Chapter 7-8 Notes

1. Ionic Bondsa. Positively charged atom stuck to a negatively charged atomb. Ionic compound

i. High melting pointsii. Brittleiii. Water solubleiv.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 Diagrama. 2 Dots on each side around element symbol symbolizes valence electron count

4. Monatomic and Polyatomica. 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 typei. Fe2+, Fe3+

1. Differentiation via Roman numeralsc. Monatomic Anions

i. Mostly nonmetals1. Labeled with suffix ‘ide’

d. Polyatomic Ionsi. Made internally with covalent bonds (sulfate)

1. Bond with with ions by ionic bonds5. Binary Ions and Notations

a. Binary Ionic Compounds --> Cations then anionsb. 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 notationsii. For polyatomic ions, put subscript after parenthesis around polyatomic ion

6. Covalent Bonds --> shared electrons, not given, to get octet satisfieda. atoms grouped by covalent bonds are called moleculesb. Molecular formula (like empirical, but not clean ratios)

i. Different combinations of same ratio of elements gives different elements1. C6H12O6 --> glucose, C3H6O3 --> lactic acid

c. Structural formula is usedi. Many different kinds, common is Lewis Structure

Page 8: Chemistry Honors Notes

7. Lewis Structure formulaa. Like Lewis Dot, but all elements together in molecule represented

i. Shared electrons circled8. Single, Double, Triple Bonds

a. Single --> One shared pair between two elementsb. Double --> 2 pairs sharedc. Triple --> 3 Pairs Sharedd. No Quadruple, because full 8 electrons would be shared

9. Dash Formula --> Dash instead of two dots for shared electrons10. 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 changec. If equal electronegativity, then nonpolar

i. If electronegative atom at center, it is also nonpolard. Difference in electronegativities

i. Less than 0.4 --> nonpolar covalentii. 0.4 --> 2 --> polar covalentiii. Greater than 2 --> ionic

12. Naming Compoundsa. Named by bonds and atomsb. Ionic Compounds

i. cation first [as in chemical formula])ii. To balance the compound (for metal) use Roman Numerals

c. Hydratesi. Absorb water into their solid structures

1. Anhydrous substances are hydrates without their waterii. Name is name of anhydrous substance and then prefix-hydrate

1. pentahydrate, octahydrate, trihydrateiii. In chemical notation, it is times sign and then hydrate with number before it

d. Molecular Compoundsi. 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 theree. Acids

i. Hydrogen + Anion1. 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 acid13. Structural Formula shows bonds, not bond shape

a. Ball and Stick model

Page 9: Chemistry Honors Notes

i. Ball is nucleus, sticks are one bond or two bonds (straight or curvy)ii. Used to give 3D representation of atom

b. All shapes are symmetrical, electrons repulse each other14. VSEPR

a. Electrons repulse each other because of chargei. Electrons placed as far from each other as possible

15. Molecular Shape and Categoriesa. Bond Angleb. Shapec. Bond Polarity

16. Molecular Shapesa. Linear

i. Molecules arranged in a straight line1. All diatomic molecules

ii. 180 degrees bond angleb. Trigonal Planar

i. Flat and triangular shapeii. Two bonds are single, one is doubleiii. 120 degrees

c. Tetrahedrali. 109.5 degreesii. One central atom and four around it (normally) and no lone pairs

d. Trigonal Pyramidali. Three pairs and one lone pairii. 107 degreesiii. Lone pair takes up more space because has higher repulsion

e. Benti. 105 degreesii. Looks like linear, but 2 unshared pairs push it further, making it bent to a side

17. Hybrid Orbitalsa. When atoms come together, their orbitals morph

i. Linear/Triple Bonds --> sp1. When atom has two sets of atoms around it2. Triple has it because at most triple bonds can exist, so only 2 atoms bonding

ii. Trigonal Planar --> sp21. When atom has 3 sets around it (triangle)

iii. Tetrahedral/Pyramidal/Bent --> sp31. When 4 sets around it (tetrahedron)

18. Bond Lengtha. As you move down, bonds are longer, move across, bonds are shorterb. Multiple are shorter than singular

19. Polarity

Page 10: Chemistry Honors Notes

a. Shape and bonds determine polarity of a moleculeb. All nonpolar charges means nonpolarc. Formaldehyde

i. H2OC (all electrons go to O, so one end is charged, other isn’t)d. Carbon Dioxide

i. Carbon Oxygen bond is polarii. Bonds face in opposite direction, polarity is cancelled outiii. Since no polar and nonpolar end, cannot easily attract to another, and is a gas at

room temperaturee. 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

molecules1. Hydrogen’s positive center attracts another oxygen’s negative, and chain is made

f. Larger Moleculesi. Polarity determines shape

1. Similar polarities attract one another2. Shape can be determined by polarity

Page 11: Chemistry Honors Notes

Chapter 12-13

1. All reactions need energya. Breaking bonds requires, making releases

2. Thermochemistry is study of heat in chemistry (part of thermodynamics)3. Exothermic vs Endothermic

a. Products - Reactantsb. Exothermic --> more heat released than absorbedc. Endothermic --> more energy absorbed than released

4. Total Energy --> kinetic and potential energy of particlesa. Enthalpy is a little more

i. Difference in enthalpy of products and reactants is energyb. Positive for endothermic, negative for exothermicc. 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 ac-

countd. Heat and moles are proportional

5. Hess’ Lawa. Net enthalpy change = sum of individual steps needed in reactionb. If coefficients multiplied by factors, enthalpy multiplied too (heat and moles are propor-

tional)c. If arrow is changed, sign is changed (endothermic becomes exothermic)

6. Calorimetrya. Heat flow and its measurement

i. Delta H by delta H of surroundings (qrxn vs qsurr)b. Heat capacity --> energy to raise object by 1degreec. Specific heat --> raise by 1g of object by 1 degreed. Calorimeter --> insulated cup with known amount of water and thermometer

i. If water heats up, exothermic, vice versaii. Measured in J or kJ

7. Caloric Theory --> heat moved from hot things to cold things8. Kinetic Theory --> heat is transfer of kinetic energy from hot to cold9. Gases have uniform traits

a. 22.4L at STPb. Mostly made of molecules (except for noble gases)c. Have mass and can be compressedd. 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 containera. 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

Page 12: Chemistry Honors Notes

11. Kinetic Molecular Modela. Have mass, in constant motion, lot of space between them, collisions are perfectly elastic,

heat makes them move faster12. 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 Tempc. Avogadro’s Law: All gases at same temp and pressure occupy same volumed. Dalton: Sum of individual partial pressures is same as total pressures

14. Ideal Gas Lawa. PV= nRT

i. P is pressure, V is volume, n is moles, R = 0.0821 atm-L/molK, T is Kelvin/Celciusii. 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 gasa. Helium safer than hydrogen, easier than hot air

16. Law Math Conversions and Mnemonic Devicesa . P T Vb. 1 atm = 101.3 kPa = 760.0 mmHgc. For R, convert all units to correspond with R (atm L / moles K) (Pa m^3/mol-K) (J/mol K)

Page 13: Chemistry Honors Notes

Chapter 14-15

1. Liquids vs Solids vs Gasesa. Gases take up whole container and match shape, liquids only match shape, solids retain

shape sizeb. Liquids barely expand upon heating and barely compress (unlike gases) and move and dif-

fuse slower than gasesc. Solids have highest attraction forces and barely moved. 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 typesa. Ionic --> holds two different elements together (solids at room temp), strong, pulled apart

by water when they dissolveb. Metallic --> holds two metals together (solids at room temp), strongc. 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. Liquidsa. Viscosity is intermolecular force dependent function

i. Resistance to friction, stronger intermolecular forces give higher viscosityii. 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. Solidsa. Crystalline -->ordered unit cells (repeating pieces in larger structure) if solid has water, it’s

hydrateb. 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 viscositiesc. Bonds in solids depend on intermolecular forces (hardness, melting points, electrical con-

ducting)i. Metallic solids --> all ions together in sea of electrons, good conductors, malleable and

ductileii. Molecular solids --> weaker, cannot conduct, higher melting point --> higher inter-

moleculariii. Ionic --> strong forces of attraction, brittle, bad conductors (unless liquids), most

stableiv. Covalent network --> very strong, made with covalent bonds, aren’t molecules

Page 14: Chemistry Honors Notes

6. Changes of state and energy involveda. (Gas --> liquid --> solid) --> exothermic, <--- endothermicb. Vaporization --> making liquid into gas, is endothermic

i. Anything over freezing is vaporizing, certain water particles are leaving surface to be-come gas

ii. Volatile --> vaporizes fastiii. 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 forcesii. In closed atmosphere, as more evaporate, more condense, until equilibrium is reached

d. Vapor pressure --> pressure of vaporized molecules, at equilibrium pressure is constanti. Any substance’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 Depositioni. Sublimation --> solid to gas, opposite is depositionii. 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 changesii. For water, heat rises, and at 100 C, heat is kept on added until water completely evapo-

rates, then temperature of steam rises until a certain pointiii. Phase diagrams --> specific temperatures and pressures where two phases can oc-

curiv. 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 phaseb. Solid solutions are alloys, gaseous is air, liquid is just liquid (aqueous means solvent is wa-

ter)i. If aqueous solutions have ions that can carry electricity, electrolytes, otherwise non-

electrolytes8. 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

solutiond. 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 Forcesa. Intermolecular forces --> between solute and solvent

Page 15: Chemistry Honors Notes

b. Solvation --> interaction between solutes and solvents (hydration if solvent is water)i. Ex: water molecules pull Na and Cl apart in NaCl in waterii. In solution, solutes and solvents are broken and intermingled (breaking bonds is en-

dothermic, 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 pressureb. 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 themi. If temperature drops after adding solute, heating will help, if stays same heating won’t

help, if increases, heating will decrease itd. 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 ite. 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 identitya. Vapor Pressure Reduction --> nonvolatile solute makes less surface contact for solvent

molecules to surface, less of them become gas (vapor pressure is proportional to concen-tration <-- 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 i-

factorc. Freezing Point Depression --> lowering freezing point, same nonvolatile rules

i. Also proportional to molarity and given constantd. Osmotic Pressure --> (solvent passes, solute does not)

i. After certain difference on both sides of semi-permeable membrane, osmosis stopsii. Without osmotic pressure, solution is isotoniciii. 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

Page 16: Chemistry Honors Notes

Chapter 16-17

1. Equilibrium --> two opposing processes occurring at the same ratea. Reversible reactions --> equilibrium where products become reactants while reactants be-

come productsi. Has to be reversible on its ownii. Forward reaction proceeds to the right, backwards proceeds to the left

b. Reaction rate depends on concentrationi. 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 Constanta. Law of Mass Action --> Explains relative concentration of products and reactants in terms

of equilibrium constantb. Equilibrium constant known as Keq --> concentration of products raised to coefficients

over concentrations of reactants raised to coefficientsi. Concentration --> molarity

c. Law of Chemical Equilibrium --> every solution has a Keq (dependent on molarity) that varies by temperaturei. 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 equilibriumi. This is the equilibrium position of the solution

e. Homogenous and Heterogenous Equilibriai. If it is homogenous, then does not change state, in heterogenous, changes stateii. If it changes to solid, omit concentration of solids in Keq

f. Reaction Quotient --> uses specific data to see if reaction is at equilibriumi. If Q> Keq, then too many products, proceed to the leftii. If Q<Keq, then too many reactants, proceed to the right

3. Le Chatelier’s Principlea. Equilibrium systems can be altered at or nearing equilibriumb. If they are altered, the system changes to counter the changec. Concentration change --> if product or reactant is added, equilibrium will shift to other

side to make more of other substance out of added substanced. 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

Page 17: Chemistry Honors Notes

i. Manipulating Le Chatelier’s Principle to make more ammonia by manipulating pres-sure and temperature

4. Dissolution and Precipitationa. In a solution, ionic solids separate into +/- ions

i. Solution has to be a polar liquidb. Process of ionic solids becoming ions is called dissolutionc. If one ion collides with the crystal, becomes part of the crystal

i. If one ion touches another ion, both neutralize and become a solidd. Ions becoming solids again from ionic form is called precipitation

i. Dissolution and precipitation are opposite systemse. 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 timeii. When the rate of dissolution and precipitation is same, you have solubility equilibrium

5. Ksp --> Ksp is solubility product, is like Keq except no denominatora. No denominator because reactant is an ionic solid

i. So it is product of products’ concentrations (at equilibrium) to powers of coefficientsii. 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 pptii. Ppt will keep on forming until Q=Ksp

6. Reactions that make precipitate are precipitate reactionsa. All are double replacement reactionsb. 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

predictd. All are double replacement because formation of ppt drives it further

i. Solids have less energy than liquids, so tend to be more stableii. Other driving forces are creation of water and gases

7. Ionic Equationsa. Complete ionic equations have everything (all the spectator ions, those who do not un-

dergo a change)i. When writing complete ionic equations, do not break up solids

b. Net ionic equations exclude spectator ions8. Common Ion Effect

a. If two ionic substances sharing a common ion are mixed (one already in a saturated solu-tion), increase of the added ion will result in ppt formingi. You are increasing the concentration of one ion past saturation (equilibrium), so ppt

must formb. Also decreases solubility for newer ions who are introduced

Page 18: Chemistry Honors Notes

Chapter 18-19

1. Properties of acids and basesa. Taste --> acids are sour and bases taste bitterb. Touch --> acids feel like water, bases feel slipperyc. Reactions with metals --> acids react with mg, zinc, iron and aluminumd. Conductivity --> both are electrolytes (ionize in water), both good conductorse. Indicators --> change color in contact with them (ex: phenolphthalein, litmus paper)

2. Neutralization --> when acids and bases meet, cancel each other’s propertiesa. In proper ratio (MaVa=MbVb), no distinctive properties remainb. Products are ionic compound (salt)

3. Definitions of acids and basesa. Arrhenius --> (properties based on when acids/bases make contact with water)

i. Acids release H+ ions (proton), bases release OH- ionsii. Acids + metals make H2 gas

b. Bronsted Lowry definitioni. 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, etcc. 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+ ions4. Strengths of acids and bases

a. Strong acid readily makes H+ ions in wateri. Indicated by ‘-->’ in chemical reactions, while weak ones have equilibrium arrows

b. Strong bases readily accept H+ ions in waterc. Strengths of acids/bases are inverse to their conjugates (strong acid=weak conjugate base)

5. Quantitative Formulas for Acids and Basesa. HA + H20 --> H30+ + A-b. 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 basica. Strong acid + strong base --> neutral saltb. Strong acid + weak base --> acidic saltc. Weak acid + strong base --> basic saltd. Weak acid + weak base --> answers may vary

7. Naming acids

Page 19: Chemistry Honors Notes

a. Acids = Binary acids --> hydrogen + element, Oxy acids --> hydrogen + oxygen + ele-ment, Carboxylic --> has carbon in it

b. Bases = Bronsted Lowry always has one unshared pair of electronsi. 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 ita. However, very little of of it is there

i. At equilibrium (25 C), there are exactly 1.0 * 10^-7 M of bothb. Kw --> 1.0 * 10^-14

i. This means that [OH-]x[H30+] = 1.0*10^-14 every timec. pH is -log[H3O+]

i. More H3O+ means low pH, low pH means acidic, high pH means basicii. Increase of pH by 1 is tenfold change

d. pH found by indicators9. 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 bufferi. 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) --> buf-

fer capacity10. Titration --> controlled neutralization reaction that lets you know concentration of acid/

basea. Measure of pH is not the same as measure of concentration of acidb. For titration, need a standard solution (known concentration), unknown solution, and indi-

catorc. After adding known amount of standard and a few drops of indicator, add unknown and at

one point color of solution will changei. Point of color change is called end point (where more of unknown than standard in so-

lution)d. Equivalence point --> when both are in same concentrations

i. Equivalence point is very close to end point, so end point is appx equivalence pointe. MaVa=MbVb, only unknown is Mb (moles of unknown), can be figured outf. 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

Page 20: Chemistry Honors Notes

Chapter 22-23

1. Chemical kinetics --> speed at which reactions occura. Chemical reactions take time to occurb. Reaction rate is the change in concentration of reactants and products over a period of time

i. Expressed in M/sc. Average rate of reaction --> (delta(reactants or products))/delta(time)

i. Rate changes, so it makes sense to find average rate2. Reaction mechanisms --> series of steps that lead from reactants to products

a. Each step is an elementary step to the entire processb. Elementary steps must add up and (through canceling) equal the original reactionc. Some substances are made in elementary steps but cancel out

i. These are intermediary productsd. Speed of chemical reaction is dependent on slowest elementary step

i. This step is called rate-determining step3. 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 reactantsii. x and y are powers of concentrations of reactants (must be found experimentally)

1. Not always the co-effecientsiii. k is a fixed constant per reactioniv. Not all reactants are involved, those whose change in concentration do not affect

reaction are excluded4. 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 effectivei. 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 (ori-entation) and also both need to have enough energy to break and make bonds

5. Sufficient energy can be found from potential energy and kinetic energya. Potential energy is energy waiting to be released, kinetic is released energyb. Energy required to break bonds comes from kinetic energy

i. Kinetic energy is determined via mass and velocity of particle6. Activation energy

a. Energy of reactants must be raised to start reactioni. Difference between max energy of reactants and energy of products is activation en-

ergyb. Activation complex --> In transition state (brief moment where substance is neither prod-

uct nor reactant) the substance is called activation complex

Page 21: Chemistry Honors Notes

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 reactants7. Five factors affect reaction rates: nature of reactants, temperature, concentration, surface

area, and catalystsa. Nature of reactants --> reactions that require slight rearrangement of reactants are faster

i. Reactions that need many covalent bonds to be broken are slowii. Reactions with reactants in different states are heterogenous (vice versa for homoge-

nous)1. Reaction between two gases will happen faster than two liquids or solids

b. Temperature --> higher temperatures lead to faster reactionsi. Molecules move faster with higher temperaturesii. Common rule of thumb --> increasing by 10 Celsius doubles or triples reaction rate

c. Concentration --> more particles to collide means faster reactiond. Surface Area --> Larger surface area means more space for reactions to happene. Catalysts --> Catalyst is a substance that increases the rate of reaction without being used

up i. Catalysts in the body are called enzymesii. Catalysts are opposites of intermediary products, are consumed first and then remade

lateriii. Catalysts lower activation energy, so more molecules can create productsiv. Substances that reduce reaction rates are called inhibitors