Chem Final

Chapter 2: Matter & Change -Matter (anything that has mass and takes up space) -Substance -Element

-Compound

-Mixture (blend of 2+substances) -Homogeneous -Heterogeneous (air, salt water) (sand, paint, salad dressing, big particle size, thick/cloudy)

physical property: conditions of a substance that can be observed without changing the substance’s composition (green, rough) physical change: only affects physical properties (cutting, bending) chemical property: ability of a substance to undergo a chemical reaction and form a new substance (react with O2 when heated, decomposing) chemical change: always produces one or more new substances -energy is absorbed or released -change in color or odor -production of a gas (bubbles) -formation of a solid from a liquid -difficult to reverse Law of conservation of mass: mass is neither created nor destroyed in any physical change or chemical reaction gas: no definite shape; no definite volume solid: definite shape and definite volume liquid: no definite shape; definite volume filtration: separates hetero; mixture poured through filter paper (coffeegrounds and water) homogenous separation: chromatography: -molecules are separated due to the attraction that they have for the paper -mobile phase: solvent- substance that moves -stationary phase: filter paper- substance that stays still crystallization: evaporate water from a solution; crystals are left behind; produce very pure solids distillation: liquid is boiled to produce a vapor and then is condensed to a liquid *review names of elements*

Chapter 3: Scientific Measurement PREFIXES: mega M 1 000 000 kilo k 1 000 deci d .1 centi c .01 milli m .001 micro µ .000 001 nano n .000 000 001 pico p .000 000 000 001

UNITS: length meter m mass kilogram kg temperature kelvin K time second s amount of substance mol m DERIVED UNITS: volume cubic meter m^3 density g/cm^3 g/mL pressure pascal Pa energy joule J

Significant Figures: all non-zero digits, zeroes in between 2 digits, right of decimal number, and end of number. (unlimited if counting # or defined quantity) + and - : round to same # decimal places as number with least # of decimal places x and / : round to same # of sigfigs as measurement with least # of sigfigs scientific notation: only include number before x sign % error: (experimental-accepted)/accepted x 100 volume: the amount of space an object takes up mass: the amount of matter in an object density: ratio of mass of object to volume; d= m/v (g/mL g/cm^3) lab equipment: see attached worksheet K=C+273 Kelvin- zero point at 0 K, absolute zero Celsius- freezing, melting 0, boiling 100 Chapter 4: Problem Solving in Chemistry -use conversion factors (ratios) for dimensional analysis Chapter 5: Atomic Structure & the Periodic Table -Dalton’s Atomic Theory -all elements are made up indivisible particles call atoms -atoms of the same element are exactly alike -compounds are formed by joining 2+ elements in whole # ratios -chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element cannot be changed to atoms of another element (1, 2, 3b WRONG) -Thomson -discovered electron using cathode rays (- charged particles he called electrons) -atom is a ball of + charge containing electron

-Rutherford -shot alpha particles at thin sheet of gold foil and proposed that atom is mostly empty space with all of the + charge in center (nucleus) -Crookes -invented the Crookes tubes -Bohr -electrons are in energy levels that orbit the nucleus; lowest energy level is closest to the nucleus mass # (# of protons and neutrons) atomic # (# of protons) SYMBOL isotopes: atoms that have the same number of protons, and different neutrons (Carbon-12 vs. Carbon-13) average atomic mass: the average mass of a mixture of isotopes for an element (atomic weight) (relative abundance)(AMU) + (relative abundance 2)(AMU 2) = average atomic mass periodic table: -left of staircase = metals (HYDROGEN IS A NONMETAL) -right of staircase = nonmetals (ALUMINUM IS A METAL) -at staircase = metalloids (have some properties of both) Chapter 6: Chemical Names and Formulas formula units: used to represent ionic compounds, lowest whole # ratio of ions in the compound to make it electrically neutral diatomic elements: N2, O2, F2, Cl2, H2, I2, Br2 polyatomic ions: OH-(hydroxide), NO3- (nitrate), ClO3- (chlorate), CO32- (carbonate), SO42(sulfate), PO43-(phosphate), NH4+ (ammonium) transition metals: copper, iron, mercury, lead, tin, chromium, manganese, cobalt Ag+, Cd2+, Zn2+ molecular compounds: *we use the stock system -molecule: smallest electrically neutral unit of a substance -made up of 2+ atoms that exist as 1 unit -compounds composed of molecules are called molecular/covalent compounds -most have low melting and boiling points ` -most exist as liquids or gases at room temp. -2 NON METALS -FORMULA: don’t need to be reduced like ionic compounds; prefixes EXCEPT when one of 1st element + ide (mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca) -REMEMBER PREFIXES

ionic compounds: -compounds composed of cations (+ charge, formed from metallic elements) and anions (- charge, “ide” ending, formed from nonmetallic elements) -high melting and boiling points -solids at room temperature -FORMULA- symbols (metal first), charge, subscripts to compound is neutral -NAMING- name of metal + name of nonmetal + ide polyatomic ions, etc: NH4+ (Ammonium) OH- (hydroxide) C2H3O2- (Acetate) NO3(Nitrate) CO32- (Carbonate) SO42- (Sulfate) PO43- (Phosphate) don’t change: Ag+ Cd2+ Zn2+ Chapter 7: Chemical Quantities molar mass: the mass in g of 1 mole of a substance (use masses on PT; O2=32!!) Avogadro’s number: # of representative particles in 1 mole, 6.02 x 1023 , MULTIPLY BY # OF ATOMS IF FINDING ATOMS VOLUME OF A GAS AT STP (0C, (273K) 1 atm. (101.3 kPa) * 22.4 L of 22.4 dm3

/ 22.4 L or 22.4 dm3 MOLES * 6.02 * 1023 (* atoms)

/ molar mass MASS

*molar mass /6.02*1023 (/ atoms)

PARTICLES

percent composition: % mass = (mass element/mass compound) * 100 empirical formulas: -gives lowest whole number ratio of the atoms in a compound (can be same as molecular formula, but not always – H2O2 vs. HO) -change to moles, divide by smaller mole, determine ratio for subscripts molecular formulas: gives you mass; find empirical and increase in ratios to reach mass Chapter 8: Chemical Reactions -ALWAYS CHECK TO MAKE SURE EQUATION IS BALANCED -REMEMBER DIATOMIC ELEMENTS AND MAKE COMPOUNDS NEUTRAL predicting products: always put + with – when replacing synthesis: several elements compounds -> one product decomposition: one reactant -> several products/elements single replacement: element + compound -> element + compound (kicks out + charge) REMEMBER REACTIVITY -left more reactive then right on PT (mostly) -nonmetals group 17: activity decreases as you go down double replacement: compound + compound -> compound + compound combustion: reactants- C, H, (O) compound + O2; products- CO2 + H2O

Chapter 9: Stoichiometry -mole conversion problems (ALWAYS BALANCE FIRST!) -mole ratios as conversion factors; coefficients tell you the relative # of moles limiting reactant: -limiting reactant runs out first; excess we have extra of -convert to moles, pick one reactant and convert it to moles of the other (DO THIS RIGHT), > previous = limiting, < previous = excess, moles of limiting to moles of product using mole ratio, (moles -> grams) % yield = (actual yield/theoretical yield) * 100 -actual- your lab data -theoretical- if everything went perfectly (find using stoichiometry) Chapter 10: Phases of Matter manometer: used to measure the vapor pressure of a liquid; Hg levels higher on atmosphere side, then liquid has a higher pressure (looks like letter N) barometer: used to measure atmospheric pressure boiling point: the temperature at which the vapor pressure of the liquid is just equal to the external pressure (the boiling point of water at STP is 100C) melting point: the temperature at which a solid turns into a liquid freezing point: the temperature at which a liquid turns into a solid (reverse of melting) kinetic energy: the energy an object has because of its motion (yo-yo going up) potential energy: energy stored in an object (rubber band; yo-yo going down) Kinetic Theory of Gases: 1) gases are mostly empty space. gas particles have negligible volumes. no forces of attraction or repulsion between gas particles 2) gas particles are in constant motion 3) collision between gas particles are perfectly elastic (no loss of kinetic energy) Chapter 11: Thermochemistry heat: the energy transferred from 1 object to another due to a temperature difference temperature: measure of the average kinetic energy; it determines the direction of heat flow (measured by thermometer) enthalpy: the heat content of a system at constant pressure specific heat: (C) the quantity of heat, in joules or calories, required to raise the temperature of 1 gram of a substance 1C change in enthalpy: q = m * C * T

(g and C) (T = final-initial) (q = heat)

heat of fusion: is the heat absorbed by 1 mol of a substance when it melts heat of vaporization: is the heat needed to vaporize 1 mol of a substance -temperature remains constant as a substance melts or boils!

heat of solidification: opposite of fusion heat of condensation: opposite of vaporization endothermic process: system gains heat as its surroundings cool down, change in heat is positive exothermic process: system loses heat as its surroundings heat up, change in heat is negative DON’T FORGET “–“ SIGNS WHEN LOSES HEAT! Chapter 12: Gas Laws Boyle’s Law: P1V1=P2V2 (inversely proportional to constant T) Charles’ Law: V1/T1 = V2/T2 (directly proportional to constant P) Combined Gas Law: (P1V1)/T1 = (P2V2)/T2 Gay-Lussac’ Law: P1/T1 = P2/T2 (directly proportional to constant V) Ideal Gas Law: PV = nRT MUST be in L or dm3 (1L = 1 dm3) -V = L or mL -P = atm or kPa

-T = K -n = # OF MOLES!

-R = 8.31 (kPa), .0821 (atm) (given)

Dalton’s Law of Partial Pressures: the total pressure in a mixture of gases equals the sum of the pressures of the individual gases Graham’s Law: -diffusion – is the tendency of molecules to move toward an area of lower concentration -effusion- is the process of which a gas escapes through a tiny hole in a container rate1 / rate 2 = √ molar mass 2 / molar mass 1 Ideal vs. Real Gases real gases… -take up space (volume); -have attractive and repulsive forces; -have inelastic collisions Chapter 13: Electrons in Atoms -energy increases as level goes out; electrons in same sublevel have the same energy -2n2 = max. # of electrons in that energy level -s=1; p=3; d=5, f=7 -1, 2, 2, 3, 3, 4, 3, 4, 5, 4, 5…. orbitals: -regions where there is a high probability of finding an electron -each orbital can hold up to 2 electrons (1s, 2s, 2p…) valence electron: electrons in outermost energy level (watch # bc out of order!!) -important bc they do all the bonding, etc.; elements in same column have same ending -noble gases very stable bc outer shell is full

electron configurations: -determine total number of electrons, then fill sublevels until you run out of electrons 1) aufbau principle: electrons enter orbitals of lowest energy first 2) Pauli Exclusion principle: orbitals can hold at most 2 electrons 3) hund’s rule: before pairing of electrons occurs, only 1 electrons occupies orbitals of equal energy Chapter 14: Periodic Table -periods are horizontal rows; groups are the vertical columns Families: 1. Alkali metals (group 1, not H) (unstable in air or water; oxidize rapidly) 2. Alkaline earth metals (group 2) (can be found in earth’s crust) 3. Transition metals (d-block) (malleable and ductile) 4. inner-transition metals (f-block) (top reactive, bottom radioactive) 5. metalloids (staircase, not Al-metal) (some props. of non metals and some of metals) 6. halogens (group 17) (can be all states of matter, diatomic, highly reactive) 7. noble gases (group 18) (odorless, colorless) Trends: atomic radius: ½ the distance between the nuclei of the same atoms due to shielding effect: electrons in outer levels are less attracted to nucleus since inner electrons are blocking them ionization energy: energy needed to remove an electron from outer energy level of an atom (high-electrons held tightly; low-held loosely) electronegativity: ability of an atom to attract a shared electron to itself (values 0-4.0) A.R. inc. I.E. dec. E.N. dec.

A.R. dec. I.E. inc E.N. inc

Chapter 15: Ionic Bonding # of valence electrons is last number in group number octet rule: when forming cpds, try to achieve 8 valence electrons Ionic bonds: TRANSFER between cations and anions; write formula first, then draw each atom separately with full electrons on – and none on + Metallic bonds: the attractive forces that hold metals together; free floating valence electrons are attracted to + metal ions; gives metals their properties

Chapter 16: Covalent Bonding Lewis dot structures: valence electrons as dots covalent bonds: electrons are SHARED with single double and triple bonds in nonmetals -count up valence electrons for each atom, write symbols, dashes for shared bonds, add dots exceptions: H only needs 2 valence; some expand their octet like P and S; B only needs 6 molecular shape: based on VSEPR Theory- pairs of electrons want to repel each other *any 2 atom molecule is linear Total Pairs 2 3 4

5 6

Shared Pairs

Lone Pairs 2 3 2 4 3 2 5 6

0 0 1 0 1 2 0 0

Shape linear trigonal planar bent (angular) tetrahedral pyramidal bent (angular) trigonal bipyramidal octahedral

Bond Angle 180° 120° < 120° 109.5° <109.5° <109.5° 90°, 120° 90°

polarity: covalent bonds can be… polar- unequal sharing of electrons -slightly – end and slightly + end (dipole) nonpolar- equal sharing of electrons (generally the same atom) -ionic have EN diff > 1.8 -use EN difference to determine bond type -polar molecule has a slightly + and – ends too (all non polar bonds = non polar molecule; polar bonds arranged evenly = non polar molecule-linear, tri.plan., and tetrah.) forces of attraction between molecules: -intermolecular forces- determine whether a molecular cpd is a solid, liquid or gas -weaker than ionic and covalent bonds -Van der Waals: weakest, hold molecules together, two types: dispersion forces and dipole interactions -hydrogen bonding: formed between an H atom and an atom w. high EN in another molecule; strongest intermolecular forces (H2O forms H bonds- gives it its properties) Chapter 17: Water & Aqueous Systems properties of water: -can dissolve many substances -surface tension: inward force or pull that tends to minimize the surface area of a liquid -liquid @ room temp. -melting pt 0C ; boiling pt 100C (high) -shape: bent polarity: polar -density @ room temp: 1 g/mL -specific heat capacity: 4.18 J/ g * C

aqueous solutions: water samples containing dissolved substances -solute- substance that gets dissolved -solvent- dissolving medium (water in aqueous solution) -“like dissolves like” – for polar/nonpolar electrolyte: substance whose aqueous solution forms ions-ionic cpds ;conducts electricity nonelectrolyte: does not form ions in aqueous solution (molecular cpds); poor conductors heterogeneous aqueous systems: -suspensions: particles settle out upon standing; larger particle size that can be separated w. filtration -colloids: particle size between solution and suspension; no filtration or settling out (milk, blood, paint) Chapter 18: Solutions hydrates: compounds that have water molecules weakly bonded to them in crystals (use a prefix for water and a dot to connect) (solve: (mol of water) / (mol of anhydrous salt) molarity: concentration of a solution. = (moles of solute) / (liters of solution) molality: concentration of a solution. = (moles of solute) / kg of solvent) colligative properties: (only depends on # of solute particles) -freezing point depression: (ΔTf) difference in temp. between freezing pt of a solution and that of pure solvent. ΔTf = Kf (molal freezing pt constant) * m * (# of particles) -boiling point elevation: (ΔTb) difference in temp. between boiling pt of a solution and that of pure solvent ΔTb = Kb (molal boiling pt constant) * m * (# of particles) factors affecting solubility: 1. temperature – increase temp, dissolve faster 2. size – smaller dissolve faster 3. stirring - faster 4. already dissolved solute - slower solubility curves: show how much solute will dissolve in a given amount of solvent over a range of temps. -unsaturated- amount of solute is much less than solubility @ the given temp. -saturated – solution contains all solute it can hold @ given temp. -supersaturated- contains more solute than it should theoretically hold @ given temp. dilutions: add more solvent to a solution; moles of solute stay the same, concentration changes. M1 * V1 = M2 * V2 (molarity and volume) Chapter 19: Rates and Equilibrium collision theory: rates measure the speed of any change that occurs over an interval of time; particles must collide in order for reactions to occur -two factors: frequency and effectiveness of collisions *a catalyst would lower the hills

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-Le Chatelier’s Principle: if a stress is applied to a system at equilibrium, a shift occurs to relieve the stress I. concentration- increase left, shifts right ( makes more products); increase right, shifts left (make more reactants) II. temperatureexo, increase = left exo, decrease = right endo, increase = right endo, decrease = left III. pressure- *only effects substances in gas phase; pressure exerted by gas is directly proportional to the # of molecules; inc pressure shifts toward fewer molecules, dec pressure shifts toward more molecules *catalyst dos NOT affect if aA + bB = cC + dD… Keq = [C]c * [D]d / [A]a * [B]b >1 products favored; <1 reactants favored *use ICE chart if needed *only gases

enthalpy: total heat content of system; exo ΔH is -, endo ΔH is + ; ΔH = sum of products – sum of reactants entropy: ΔS = S(products) – S(reactants) + : entropy increases, favorable - : entropy decreases, unfavorable free energy: ΔG = ΔH – T ΔS (T is temp. in Kelvins, *have to change ΔS to kJ) ΔG = G(products) – G(reactants) + : nonspontaneous (will not occur) - : spontaneous (occurs naturally) rates of reactions (Kinetics): rate = K [A] -order of reaction = power to which concentration of reactants must be raised in rate law -find order by seeing how much rate inc when inc concentration (nothing-0th order, double- 1st order, quadruples, 2nd order….) -overall order- add superscripts together Chapter 20: Acids and Bases acids: sour taste; aqueous solution of acids are electrolytes; red indicator; < 7 pH produces hydrogen ions when dissolved in water HX -ends in –ide = hydro_____ic acid (no O’s present) -ends in -ite = _____ous acid -ends in –ate = _____ic acid

bases: when aqueous- taste bitter, feel slippery and are electrolytes; blue indicator; >7 pH produces OH- when dissolved in water name as normal ionic compound pH = -log[H+] pOH = -log[OH-] pH + pOH = 14

[H+] = 10 -pH [OH-] = 10 –pOH

Kw = [H+] * [OH-] = 1.0 * 10-14 -strong acids completely dissociate in water; weak acids only partially ionize in water Ka = [C]c * [D]d / [A]a * [B]b *use ICE chart small = weak; large = strong Chapter 21: Neutralization neutralization reaction: acid + base -> salt + water titrations: used to determine concentration of a solution standard solution- the solution of known concentration indicator- added to signal when neutralization has occurred end point/equivalence point- when the neutralization has been reached N1V1 = N2V2 V=mL or L N=M, 2M, 3M… based on # of acid/base ions *see attached sheet for solubility curves! salts: weak acid neutralized by strong base = basic salt solutions strong acids neutralized by weak bases = acidic salt solutions strong acids neutralized by strong bases = neutral salt solutions Ksp = (same as other K equations – do not use solids!) the larger the Ksp, the more soluble the ionic cpd is in water Chapter 22: Redox LEO the lion says GER -Lose Electrons Oxidation, Gain Electrons Reduction oxidation numbers: 1)monatomic ion = charge of ion 2) H is +1 except in metal hydrides it is -1 3) O is -2 except peroxides it is -1 and OF2 it is +2 4) any free element is 0 5) sum of oxidation #’s in neutral molecule is 0 6) sum of oxidation #’s for polyatomic ion =’s the charge of the ion

Chapters 25 & 26: Organic Names and Formulas: -use proper punctuation - commas separate #s and hyphens separate #’s and words. attached groups go in alpha. order, but ignore prefixes. halogens first, then alkyl groups -use longest Carbon chain; a substituent comes off of chain, an alkyl group is a type of substituent with H’s and C’s -#’s say where attached group is coming off of, prefix tells how many C’s or groups (meth, eth, prop, but, pent, hex, hept, oct, non, dec) *see attached chart structural isomers: same molecular formula, but different structure and properties (CH3CH2OH and CH3OCH3) geometric isomers: cis isomer- groups are on the same side of double bond; trans isomergroups lie across double bond stereoisomers: molecules that are mirror images of each other *see attached worksheet for functional groups Chapter 28: Nuclear Chemistry half life: time it takes for ½ of atoms in sample to decay

amount remaining = original * (1/2) (time decay/half life) alpha radiation: 2 protons and 2 neutrons 42He

or

4

2

α (mass over charge)

beta radiation: -1e or -1 β gamma radiation: mass of 0 amu, no charge, type of electromagnetic radiation, similar to x-rays, but more dangerous, γ 0

0

nuclear reactions: superscripts = conversion of mass; subscripts = conversion of charge Chapter 23: Electrochemistry electrolytic cells: an electrochemical cell used to cause a chemical change through the application of electrical energy voltaic cells: an electrochemical cell used to convert chemical energy into electrical energy; the energy is produced by a spontaneous redox reaction same: electrons flow from anode (oxidation) to cathode (reduction) in the external circuit differences: voltaic cell- flow of electrons is result of spontaneous redox rxn; anode is electrolytic cell- electrons are pushed by an outside power such as battery; cathode is – (bc attached to – side of battery)