# SAT Review

Flashcard maker : Viola Marenco
 1. Mass
 amount of matter in a sample matter = anything that occupies space & has mass grams, milligrams, kilograms more mass means more atoms/molecules
 1. Volume
 how much space sometimg takes up usually liters or milliliters 1 cm3 = 1 mL measure liquid volume with measuring flask measure solid volume with displacement gas volume = volume of container (displacement with container if necessary)
 1. Density
 ratio of mass to volume of an object each substance has a specific density d = m/v density of liquids and solids as a fixed temp is constant density of gases changes
 1. Presh
 force that gas in a closed container exerts on the container walls solid/liquid is there, gas exerts presh on walls of environment and everything in it including the solid/liquid use manometer or barometer to measure (both use Hg) 760 torr = 760 mmHg = 1 atm
 1. Energy
 the ability to do work or transfer heat heat, light, kinetic, chemical bond energy (different forms) usually kinetic energy of molecules in chem–greater KE means higher temp and faster movement breaking bonds takes energy, forming them releases energy heat is the transfer of KE from one thing to another 1 cal = 4.186 J; 1000 J = 1 kJ measure it with a calorimeter
 1. Temperature and Heat
 heat is energy flow from high temp thing to low temp thing temp measures average KE of molecules in a sample if something increases in temp, that’s a heat content increase substances have different specific heats K = C + 273; 0 K = -273 C; 0 C = 273 K
 1. Specific Heat
 it’s easier to heat some things than others heat capacity is how much energy something has to absorb for some of it to be raised 1 degree Celsius specific heat is the heat capacity for 1 g of the substance q = mst
 2. Atom
 smallest particle of an element retains the chemical properties of the element
 2. Parts of an Atom
 proton, positive charge neutron, no charge nucleons, in the nucleus, protons and neutrons electrons, negative charge nucleons have mass; electrons have practically none neutrons determine the isotope
 2. Ion
 atom with unequal charges inside and outside nucleus cation, positively charged anion, negatively charged
 2. Element
 most fundamental unit of matter can’t break it down without losing its identity
 2. Table Organization
 horizontal rows are periods vertical columns are groups all elements in the same group have the same number of electrons in outer shells and share properties
 2. Atomic Number
 number of protons in the nucleus of an atom specific to the element
 2. Mass Number
 protons and neutrons have 1 amu mass sum of an atom’s proton and neutrons is the mass number isotopes have different numbers of neutrons in their nuclei doesn’t appear on the table because it varies
 2. Atomic Weight
 average mass number, so it’s based on each isotope’s natural abundance one for each element listed on the table
 3. Molecule
 units of two or more atoms held together with chemical bonds
 3. Diatomic Molecule
 molecule made of just two atoms some elements exist this way at stp: O2, I2, H2, N2, Cl2, F2, Br2 Clearly I Have NO Friends, Bro
 3. Formula Weight
 add the atomic weights of all the atoms in the molecule
 3. Empirical Formula and Molecular Formula
 empirical formula is the ratio of atoms within a molecule the molecular formula divided by the gcf of the subscripts gives the empirical formula
 3. Percent composition
 refers to percent by mass find it by calculating stuff 😛
 3. Mole
 a mole of something is 6.022 x 1023 the 6 number is called Avogadro’s number a mole of atoms make up the atomic mass
 3. Mass Composition to Empirical Formula
 imagine 100g convert percents to grams convert grams to moles find the ratio between the mole amounts
 3. Stoichiometry Problems
 how much product or reactant is produced or needed in a reaction? always start with a balanced equation remember limiting reactants convert things to moles!
 3. Entropy
 S symbol If deltaS is negative, reaction loses entropy If deltaS is positive, reaction gains entropy Universe tends towards disorder Low energy states ar emore stable than high energy states Higher S + Lower E = More Stable
 3. Enthalpy
 term refers to energy states of reactions or products symbolized with H decreasing H leads to stability H decreases: exothermic; H increases: endothermic endothermic reactions need energy input energy of products/reactions refers to the energy in the bonds
 3. Heat of Formation
 amount of heat released or absorbed when a mole of a compound is formed if Hf is negative, then exothermic, and vice versa for all elements, the heat of formation is zero heat of formation for entire reaction = sum of product Hfs minus sum of reactant Hfs Hess’s Law
 3. Hess’s Law
 If a reaction happens in more than one step, the change in enthalpy for the whole thing is the sum of the changes in enthalpy for each step Enthalpy is pathway-independent: state function
 3. Spontaneous Reaction
 happens without energy input can happen for a positive H if S is low enough can happen for a positive S if H is low enough spontaneity is determined by Gibbs free energy
 3. Gibbs Free Energy
 ΔG =  ΔH – TΔS T is in Kelvins determines if a reaction will happen spontaneously or not if ΔG < 0, then it’s spontaneous in that direction if;;G ; 0, then it’s spontaneous in the opposite direction if ΔG = 0, then there’s equilibrium
 4. Quantum Mechanics
 our current theory about how electrons and atoms work a quantum is a small unit of energy, and all energy is quantized, so all energy exists in multiples of quanta
 4. Orbitals
 electrons exist in them replaces Bohr model of orbiting describes the likelihood that an electron will be found in a particular location–a probability function orbitals have energy shells, shape, subshells any orbital holds two electrons
 4. Energy Shell
 each orbital is in an energy shell higher number energy shell greater energy farther from nucleus each energy shell has a whole number
 4. Orbital Shape and Subshell
 4 different shapes that make up subshells s has 1 orbital p has 3 orbitals d has 5 orbitals f has 7 orbitals
 4. Heisenberg Principle
 it’s impossible to know both the position and momentum of an electron at the same time
 4. Louis De Broglie
 matter has the properties of a wave electrons behave in waves like electromagnetic radiation
 4. Bohr Model
 wrong electrons orbit the nucleus in true orbits, like how planets orbit the sun
 4. Aufbau Principle
 each subshell is filled before electrons go in the next subshell Cr and Cu are exceptions (promote a 4s to a 3d) Hybrid orbitals are exceptions
 Unstable nucleus undergoes nuclear decay –> radioactive –> releases radioactivity radioactive particles alpha beta gamma as a radioactive atom decays becomes another isotope or becomes another element
 4. Geiger Counter
 4. Four Types of Radioactive Decay
 alpha decay beta decay positron emission gamma decay
 4. Alpha Decay
 emits alpha particle: 2 pros and 2 neus decreases atomic # by 2 and atomic mass by 4 sometimes symbolized as 4/2He
 4. Beta Decay
 emits a beta particle: one electron releases an electron from a neutron –> lose neutron and gain proton atomic # increases by 1 and mass stays the same
 4. Positron Emission
 releases positron: antiparticle of an electron, same magnitude but opposite charge proton becomes a neutron decreases atomic # by 1, mass stays the same
 4. Gamma Decay
 emits gamma rays when the atom has too much energy makes nucleus more stable but doesn’t do anything else
 4. Half-life
 rate of a substance’s radioactive decay takes one half-life for half of the substance to decay
 5. Horizontal and Vertical Similarities
 same row (period) –> electrons in same energy shells same column (group) –> similar chemical and physical properties
 5. Family Traits
 same column noble gases alkali metals alkaline earth metals halogens
 5. Alkali Metals
 group 1/1A 1 valence electron very reactive snihy, grayish-white melt easily lower densities than other metals
 5. Alkaline Earth Metals
 group 2/2A 2 valence electrons less reactive than alkali metals but more reactive than other metals
 5. The Active Metals
 Alkali metals Alkaline earth metals they’re really reactive…duh
 5. Halogens
 group 17/7A 7 valence electrons very reactive look different from each other
 5. Metals
 shiny and conduct heat well malleable and ductile all solid at room temp except mercury lose electrons in bonds
 5. Nonmetals
 share or gain electrons in bonds poor conductors of heat and electricity some are solid, liquid, gas at room temp
 5. Semimetals/Metalloids
 some characteristics of metals and nonmentals can either gain or lose electrons in a bond
 5. Periodic trends to know
 Ionization energy Electronegativity Atomic radius Metallic character
 increases towards cesium
 5. Ionization Energy
 energy required to remove an electron from an atom increases towards fluorine 1st, 2nd, etc.
 5. Electronegativity
 pull that a nucleus exerts on electrons of another atom increases towards fluorine
 5. Metallic character
 5. Lattice Energy
 energy required to completely separate a mole of a solid ionic compound into separate ions higer lattice energy = stronger bond
 5. Electrostatic Force
 attraction betweena positive charge and a negative charge strong
 5. Transition Metals
 middle of the periodic table harder than metals higher melting point than metals compounds with transition metals are colorful
 6. Ideal Gas Assumptions
 molecules of an ideal gas don’t attract or repel each other molecules of an ideal gas occupy zero volume
 6. Kinetic Molecular Theory
 kinetic energy of a gas molecule increases proportionally with temperature in K
 6. Network Solid
 atoms bonded by covalent bonds in a continuous network one big macromolecule no individual molecules e.g. diamonds or quartz
 6. Hydrate
 Ionic substance where water molecules bond to the ions in a fixed ratio dot H2O means it’s a hydrate % comp of water by mass in a hydrate is called water of hydration
 6. Four Types of Crystalline Solids
 Ionic Covalent network Molecular Metallic
 6. Heat of Fusion
 amount of heat it takes for a substance to go from solid to liquid
 6. Heat of Vaporization
 Amount of heat it takes for a substance to go from liquid to gas
 6. Phase change and presh
 more pressure means it’s harder to get solids to melt or liquids to vaporize more presssure does the opposite for water–easier to melt
 6. Triple Point
 a particular temperature and pressure substance can exist as a solid, liquid, or gas
 6. Vapor Presh
 created when liquids below their boiling points evaporate all liquids in a closed system exert some higher vapor presh means more volatility affected by intermolecular forces, temperature, and molecular weight boiling all of the presh above a liquid is vapor presh vapor presh = atmospheric presh
 6. Energy and Phase Change
 solid ———–> gas low PE ——–> high PE, which universe dislikes low entropy –> high entropy, which universe likes melting/boiling is spontaneous when the temperature is above the melting/boiling point
 7. Molarity
 M M = moles of solute/liters of solution
 7. Molality
 m m = moles of solute/kilograms of solvent
 7. Solubility Factors
 Solids more soluble at higher temperatures Gases less soluble at higher temperatures more soluble at higher presh think about carbonated drinks
 7. Electrolytes
 ions in solution makes the solution about to conduct electricity the solution is still neutral in ionic solutions aka elecrolytic solutions
 7. Solutes and Freezing/Boiling Points
 ΔT = kmi solute effects an increase in boiling point and decrease in freezing point k changes from solvent to solvent m = molality i = the number of particles it dissolves into per mole sucrose: i = 1 NaCl: i = 2
 7. Precipitate
 solid substance that settles out of a solution often the result of a double replacement reaction use solutbility rules
 7. Solubility Rules
 all nitrates and perchlorates are soluble all alkali metal and ammonium compounds are soluble hydroxides are insoluble except alkali metals and BaOH silver, lead, and mercury salts are insoluble except with nitrates and perchlorates
 8. Kinetics
 study of the rates of reactions
 8. Equilibrium
 the point in a chemical reaction at which concentration of reactants and products ceases to change forward and reverse reaction rates are equal
 8. Activated Complex/Transition State
 unstable place where products form need to reach activation energy before activated complex can form
 8. Factors that affect likelihood of reactions
 frequency of collision concentration of reactants surface area of reactants temperature energy of collision temperature nature of reactants catalysts
 8. Catalysts
 increase rate of reaction by lowering activation energy not consumed in a reaction don’t change the equilibrium of a reaction
 8. Equilibrium constant expression
 reaction: aA + bB <—> cC + dD Keq = ([C]c [D]d)/([A]a [B]b) solvents not included powers are coefficients from the equation Keq = products/reactants if > 1 then forward reaction favored if < 1 then reverse reaction favored
 8. Le Chatelier’s Principle
 add more on one side –> equilibrium shifts to other side Keq stays constant unless temperature changes if stress is placed on a reaction at equilibrium, then the equilibrium will shift in a direction that relieves the stress
 8. Haber Process
 makes ammonia
 8. Pressure Change Effect on Equilibrium
 only applies if some of the things in the reaction are gases reduce volume –> increase pressure –> equilibrium shifts to side with fewer moles of gas
 8. Solubility Product Constant
 Ksp Equilibrium constant between insoluble ionic solid and dissolved ions The values are usually really small Solids aren’t included in normal equilibrium expressions
 9. Autoionization
 spontaneous dissociation of water tha tmakes H+ and OH- ions reversible an equilibrium exists between H+, OH-, and H2O equilibrium expression for it is 10-14 M2 at 25 degrees C same in every aqueous solution
 9. pH
 “p” means -log if [H+] = 10-7 M, then pH = 7
 9. Different Definitions for Acid and Base
 Arrhenius Acids make H+ in aq Bases make OH- in aq Lewis acids accept electrons in aq bases donate electrons in aq Bronsted-Lowry acids donate protons (H+) bases accept protons most common definition currently
 9. Acid Dissociation Equations
 HA(aq) –> H+(aq) + A–(aq) or HA(aq) + H2O(l) –> H3O+(aq) + A–(aq)
 9. Base Reaction Equations
 A–(aq) + H+(aq) –> HA(aq) or A–(aq) + H2O(l) –> HA(aq) + OH–(aq)
 9. Amphoteric
 molecules or ions that can act as acids or bases depending on the solution most molecules are either acids or bases
 9. Strong Acids and Bases
 completely dissociate dissociation is 100% irreversible one-way arrow pH of 1.0 M strong acid is always 0 pH of 1.0 M strong base is always 14
 9. List of Strong Acids
 HCl HBr HI HNO3 H2SO4 (only the first H is strong) HClO4
 9. List of Strong Bases
 LiOH NaOH KOH other group 1 hydroxides
 9. Weak Acids and Bases
 partially and reversibly dissociate use the reversible arrow figure out if something’s acidic or basic, then it’s weak if it isn’t on the strong list remember that pH + pOH = 14 at 25 degrees C
 9. Conjugate Pairs
 molecules that are the same except one has an extra H+ the one with the extra H+ is the onjugate acid the one with fewer H+ is the conjugate base doesn’t necessarily mean that they’re an acid and a base there’s an ionic equilibrium between the conjugates sum of pKa and pKb of conjugates always = 14 at 25 C
 9. Conjugate Rules
 Conjugate acid of a strong base is neutral Conjugate base of a strong acid is neutral Conjugate acid of a weak base is an acid Conjugate base of a weak acid is a base
 9. Buffer
 minimize a change in pH when an additional acid or base goes into a solution a conjugate pair of a weak acid and a weak base so they don’t neutralize each other use Henderson-Hasselbalch equation to calculate pH
 10. Oxidation States
 Sum of oxidation states in a compound is zero Oxygen: -2 Alkali metals: +1 Alkaline Earth Metals: +2 Halogens: -1 Hydrogen: ±1
 10. Oxidizing Agent
 gets reduced causes something else to get oxidized
 10. Reducing Agent
 gets oxidized causes someone else to be reduced