MSL Review – Flashcards
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Unlock answers| Unit 1 |
Metric Conversions~ Be able to use KHDbDCM (King Henry Died Drinking Chocolate Milk) to convert (without conversion chart) Kilo(k) Hecto(h) Decka(da) -base unit- (meters, liters, seconds) Deci(d) Centi(c) Milli(m) Ex.) 1dam = 100dm Ex.) 1km = 100,000cm Sig Fig Rules~ 1. All non zeros are significant -After a non zero AND after a decimal ex.) 200.0 = 4sigfigs 3. Zeros ARE NOT significant when: -before non zeros after a decimal ex.) .002 = 1sigfig -After non zeros with no decimal ex.) 400 = 1sigfig Sig Figs in Math Problems~ -Adding & subtracting:round to least # of decimal places ex.) 1.5 + 2.250 = 3.75...round to least # of decimal places...3.8 -Multiplying & dividing:round to least # of sig figs ex.)2.40 x 1.523 = 3.6552...has 5 sig figs but the least # of sig figs in the problems is 3...3.66 Density = Mass/Volume~ -Units: g/mL or g/cm3...1mL = cm3 Matter: Classification~ 1. Element:made up of only one type of atom and cannot be broken down chemically 2. Compound:made up of more than one type of atom and can be broken down chemically 3. Mixtures:combination of elements, compounds or both; can be broken down physically -Homogeneous - same composition throughout (aka solution) -Heterogeneous - different composition throughout |
| Unit 2 |
Subatomic particles~ 1. Protons: -in a nucleus -positive charge -mass= neutron (~1 amu) -is the atomic # (defines the element!) -same # of protons as electrons in neutral atoms 2. Neutrons: -in nucleus -no charge (neutral) -mass = proton (~1amu) -number of neutrons = mass # - atomic # 3. Electrons: -Around nucleus(in energy levels) -negative charge -mass = 1/1837 of proton/neutron -same number as protons in neutral atoms -involved in bonding and determine properties of elements(what the element is going to bond with!) Ions~charged particles due to loss or gain of electrons -Cations: positively charged due to loss of electrons(metals) -Anions: negatively charged due to gain of electrons(non-metals) Isotopes~different mass # because of the different # of neutrons (# of protons remain the same!) -Element - #: the number represents the mass # of that element (i.e. Carbon-14 or C-14...the mass of carbon is 14) Average atomic mass~(of isotopes) (mass x %abudance) + (mass x %abudance) +... 100 Bohr's model~shows where electrons are (probably) found in their energy levels -Protons in center, electrons orbiting around the nucleus in energy levels(lowest energy level is closest to the nucleus) -If an electron gains energy, it has to move up energy levels - "excited state" -If an electron releases energy, it drops down energy levels - "ground state"...energy is emitteed in specific amounds called quanta(plural for quantum) which is energy that is emitted in some form of electromagnetic radiation as photons (particles that transmit light, visible and not visible) Electromagnetic Radiation~form of energy that exhibits wave-like characteristics as it travels through space -Shorter wavelength, higher frequency -Longer wavelength, lower frequency -Energy and frequency:more energy in shorter wavelenths; less energy in longer wavelenghs [image] Square of Periodic Table~ -Atomic #: Whole number; tells number of protons and electrons if atom is neutral; defines element; Symbol = Z...i.e. Z=20 the atomic number is 20, therfore the element is calcium -Atomic mass: Usually a decimal number; the mass of all subatomic particles(protons, neutrons and electrons)l Symbol = A -Symbol: Letters representing the name of the element; first letter capitalized, all others lowercased |
| Unit 3 |
Periodic Table~ -In order of increasing atomic mass:Dmitri Mendeleev -In order of increasing atomic number:Henry Moseley Period~ -Horizontal rows -Indicate principle energy levels Group(Family)~ -Vertical columns -Indicate number of valence electrons -Elements in the same family have similar properties due to the same number of valence electrons in each group *Reactivity increases as you go down for metals and increases as you go up for non-metals Periodic table classifications~ 1. Alkali metals: group 1A -highly reactive with water -one valence electron -Charge of +1 2. Alkali earth metals: group 2A -react well with halogens -two valence electrons -charge of +2 3. Transition metals: groups 3-12 -variable charges 4. Halogens: group 17 (7A) -React well with alkaline earth metals -7 valence electrons -charge of -1 5. Noble Gases: group 18 (8A) -unreactive -8 valence electrons (except Helium with only 2) -charge of 0 *Tall groups are also known as main group or representative elements (the "A" groups) Blocks~spdf -s: groups 1 ; 2 (including He), starts with energy level 1 -p: groups 13-18, starts with energy level 2 -d: groups 3-12, starts with energy level 3 -f: bottom 2 rows (lanthanides ; actinides), starts with energy level 4 Electron Configuration~ -used to show where electrons are located in energy levels and sublevels -First number represents energy level; letter represents sublevel; superscript number represents # of electrons Orbital Notation~ -Hund's rule: each orbital must be occupied by one electron before a second electron can occupy it ^v ^v ^v ^ ^ 1s2 2s2 2p6 Noble Gas Configurations~ -Use the noble gas BEFORE the element, in brackets, and continue the configuration of the element
*s sublevel: 1 orbital p sublevel: 3 orbitals d sublevel: 5 orbitals f sublevels: 7 orbitals |
| Unit 4 |
Bonding~results when atoms gain, lose or share VALENCE electrons -most atoms want 8 valence electrons, or the same number of electrons as the noble gas that is closest on the table (He is the only noble gas with only 2, since it only has an s orbital) Ionic Bonds~ 1. Between metal and non-metal (attraction between cartion and anion) 2. Electrons are TRANSFERED (metal loses & non-metal gains electrons) 3. Properties: -high melting/boiling points -soluble in water -conduct electricity in the liquid phase (electrolyte-dissolves in water and can condut electricity) -hard and brittle 4. Ionization energy - how much energy is required to remove an electron from an atom -trend: increases from left to right, bottom to top (electronegativity has the same trend) 5. Charges: -group 1A: +1 (alkali metals) -group 2A: +2 (alkali earth metals) -group 3A: +3 -group 4A: +4,-4 -group 5A: -3 -group 6A: -2 -group 7A:-1 (halogens) -group 8A: 0 (noble gases) -transition metals: variable charges 6. Naming ionic bonds: -cation name is unchanged -change anion ending to -ide if monatomic(found on periodic table), leave it the same if polyatomic -if first element is a transition metal and variable charge, put the Roman Numeral charge in parenthesis after the cation name 7. Writing formulas for ionic bonds: -cation on the left, anion on the right -write the symbol and charge of cation, symbol and charge of anion, drop charges and criss-cross numbers (should start at the top and end as subscripts at the bottom) -must always use parenthesis for polyatomic anions if the subscript is anything besides 1 -if there is a Roman Numeral in parenthesis, that indicates the charge of the metal cations *Remember - Zinc is always +2 and Silver is always +1 Covalent Bonds (aka molecules)~ 1. Between 2 non-metals 2. Electrons are SHARED 3. Bonds -single bond: 2 atoms share 2e- -double bond: 2 atoms share 4e- -triple bond: 2 atoms share 6e- 4. Properties: -low melting/boiling points -doesn't conduct electricity -solids are brittle 5. Naming covalent bonds -use prefixed to indicate # of each atom -never use the prefix mono- for the 1st element -you can use any prefix for the 1st element except mono-, and you can use mono- for the 2nd element just never the 1st *No charges or criss-crossing for covalent molecules* Metallic Bonds~ -Occurs in pure metals and metal alloys -attraction between metal ions and mobile valence electrons (electrons are free to move around the metal, causing an attraction between positive and negative charges) -the more valence electrons a metal has, the stronger the bond Electronegativity~how attracted an atom is to electrons (that are already) in a bond 1. Increases from bottom to top, left to right 2. The higher the electronegativity, the more thje atom pulls the electrons to it 3. Electronegativity difference can tell you if the bond is ionic or covalent and polar or non-polar covalent -E.D. : 0-16 = covalent »E.D. : 0-0.3 = nonpolar covalent (even sharing) »E.D. : 0-31+ = polar covalent (uneven sharing) -E.D. : 1.7+ = ionic Inter- vs. Intramolecular Forces~ -Inter: between (bonds that occur between 2 molecules) -Intra: within (bonds that occur between 2 atoms within a molecule) Types~ 1. Dipole-dipole forces -between 2 polar molecules 2. Hydrogen bonds -occur between moleucles containing H-N, H-O, or H-F bonds -hydrogen of one molecule is attracted to unshared electron pairs on another molecule -occurs because N, O and F are so electronegative, its almost as if H has lost its electron and is looking for more 3. London dispersion -only force that occurs within noble gases and non-polar molecules -electrons moving around creates instantaneous but temporary dipoles Bond Strength~ -ionic ; metalic ; covalent ; dipole-diple ; hydrogen ; london dispersion
Bond Length/Distance~the average distance between nuclei of 2 bonded atoms in a molecule. Bond length is related to bond order: when more electrons participate in bond fomration the bond will get shorter. Bond length is also inversely related to bond strength- the stronger the bond, the shorter it is. -single bonds;double bonds;triple bonds -shorter bonds = stronger bonds Atomic ; Ionic Size~ -atomic size increases from top to bottom and right to left (For top to bottom, with the increase in the number of energy levels, the size of the atom must increase. The increase in the number of energy levels in the electron cloud takes up more space. For right to left, elements on the left side of the periodic table have more protons than elements on right side. More protons make the atoms smaller becuase of their stronger pull that draws in the electrons while more electrons make the atom radii larger because of greater repulsion.) -cations get smaller, anions get larger (due to greater repulsion with more electrons) Shape - VSEPR~ 1. Valence Shell Electron Pair Repulsion 2. Electron pairs in the valence shell want to get as far away from other electron pairs as they can, so they arrange themselves around the central atom to acheive this 3. Shapes: -linear: 1 or 2 atoms bonded to central, no unshared pairs on central -bent: 2 atoms bonded to central, 2 unshared pairs on central -tetrahedral: 4 atoms bonded to cebtral, no unshared pairs on central -pyrimidal: 3 atoms bonded to central, 1 unshared pair on central -bond angles (conceptual) i.e. bond angle of a linear shape is longer than any other shape (180°) [image] Lewis Structures~Use of dot diagrams to show how atoms bond 1. Ionic: cations lose electrons so no electrons are drawn with charge, anions gain electrons so put 8 valence electrons with charge 2. Covalent: -least electronegative atom in center (although hydrogen will NEVER be central; if carbon is in the compound, it will ALWAYS be central) -all others symmetrically around central -can only bond in places where there is 1 electron!! -show single bonds with 1 line, double bonds with 2 lines and triple bonds with 3 lines |
| Unit 5 |
Chemical Chage~ -changes the identity of the substance Indicators Of a Chemical Change~(shows that a chemical change has taken place) -Formation of a gas -Color change -Formation of a precipitate -Absorption or release of heat Types of reactions~ -Single displacement: Free element + compound > different free element + new compound -Double displacement: Compound + compound > new compound + newcompound -Synthesis: one product -Decompositon: one reactant -Combustion: hydrocarbon + O2 > CO2 + H2O Predicting Products~ Balancing Equations~ 1. Law of conservation of mass- mass is neither created nor destroyed (total mass of reactants = total mass of products) 2. An unbalanced equation is called the "skeleton equation" 3. Use coefficients to get the same # of each atom on both sides 4. Can only change coefficients, NEVER subscripts -coefficients have to be in lowest whole # ratio - the most simplified 5. Can balance with polyatomic ions if ion doesn't change Activity Series~ -For single displacement reactions to see if the reaction will take place or not -One for metals and one for halogens - can replace anything below but nothing above (as you go up the activity series, the more active the element) Molar Mass~ -For a single element, equal to its atomic mass -For a compound, add all the masses of each element together (if there are multiple atoms of the element, multiply the mass by the subscript before adding) -Use when convertin grams to moles or vice versa Avogadro's Number~6.02 x 1023particles = 1 mole -use when converting atoms to moles or vice versa 22.4 L = 1 mole~ -use when converting volume of a gas at STP to moles or vice versa Percent Composition~(total mass of element/total mass of compound) x 100 -Tells what % of the compound is made up of each element -Round answer to 1 decimal Empirical vs. Molecular Formulas~ -Empirical: simplest whole # ratio of atoms in a compound; ionic compounds are always in their empirical form -Molecular:total # of atoms in a compound; covalent compounds are shown in molecular form (ex. diphosphorus hexaoxide = P2O6) Calculating Empirical Formula~ 1. Start with grams -if moles are given, skip this step -if % is given, assume its out of 100g, so change % to grams 2. Get to moles using molar mass 3. divide by the smallest mole value 4. Round to nearst whole # (as long as you're rounding up or down by no more than 0.1) -if # is too far to round (~x.1 - x.9) multiply by something to get rid of the decimal. You can multiply by 2 to get rid of .5 or by 4 to get ride of .25! Calculating Molecular Formula~Molecular mass (given in the problem) divided by empirical formula mass in grams -Should get a whole #...this is NOT the answer! -Multiply the whole # by the empirical formula (distribute to each subscript) |
| Unit 6 |
Stoichiometry~Used to calculate how much product will be produced if a given amount of reactant is added, or how much reactant needs to be added to produce a given amount of product Steps: -Balance the equation -Get to moles -Mole to mole ratio (use coefficients of starting and ending substance) -Get to desired unit (use molar mass, Avagadro's #, or 22.4L) Yields~ 1. Theoretical yield: how much product CAN be made from a given amount of reactant (calculate this with stoichiometry) 2. Actual yield: measured or experimental value; how much is ACTUALLY made from a given amount of reactant - given in problem. (The actual yield is typically a smaller amount than the theoretical yield) 3. Percent yield: ratio of actual to theoretical yeid (% of theoretical yield that is actually produced)...(actual/theoretical) x 100% |
| Unit 7 |
Phase Change~effected by temperature and pressure -higher temperature = solid>liquid>gas -The more pressure a liquid is under, the harder it is to become a gas Temperature~measure of the avg. speed of molecules Temperature Scales~Celsius and/or Kelvin are used in science - never Fahrenheit -C>K, add 273 K>C, subtract 273 -increments of C and K are the same Requirements For a Reaction~you must have particles colliding at the correct angle, and they must collide with enough energy (increasing the temp. increases the speed of the molecules and thus the energy in the collisions) Reaction Rate~speed at which reaction takes place Factors: -nature of reactants -temperature -concentration -presence of a catalyst Energy Diagrams~ -Show progression of energy throug a reaction -Activation energy: how much energy is needed to start a reaction (from the reactants to actived complex) -Activated complex: highest point on curve (where products begin to form) -Heat of Reaction (?H) - how much heat is lost or absorbed in a reaction (products - reactants) -Endothermic - heat is gained (?H is positive) -Exothermic - heat is lost (?H is negative) Catalyst~lowers activation energy (speeds up rxn) Gas Laws~shows the effects of changing pressure, volume, and temperature of a gas 1. Boyles - P1V1 = P2V2 2. Charles - V1T1 = V2T2 3. Gay Lussacs - P1/T1 = P2/T2 4. Combined - P1V1/T1 = P2V2/T2 5. Ideal - PV=nRT -P= pressure (any unit) -V= volume (C: L or mL, I: L) -n= moles (only for I) -R= gas constant (only for I; on reference table) -T= temperature (has to be in K) |
| Unit 8 |
Heat vs. Temperature~ 1. Heat- total amount of energy of each particle (a bucket of boiling water has more heat in it than a cup of boiling water even though they are at the same temperature) 2. Temperature- measure of the speed of molecules -Endothermic: absorption of heat -Exothermic: release of heat Heat Calculations~ 1. q = m?TCp (use when temperature increases) -q: heat; in calories or joules -m: mass; in grams -?T: temperature; in celsius or Kelvin (*remember, increments are the same!) -Cp: specific heat; in J/g°C, cal/g°C, J/gK, cal/gK (how much heat is required to raise 1 gram of substance by 1°C) 2. q= mHf (use when melting or freezing) -Hf: heat of fusion; how much heat is required to melt 1 gram of substance at its melting point (for water: 334 J/g) 3. q=mHv (use when vaporizing or condensing) -Hv: heat of vaporization; how much heat is required to vaporize 1 gram of substance at its boiling point (for water; 2260 J/g) 4. Melting: solid to liquid 5. Freezing: liquid to solid 6. Vaporizing: liquid to gas 7. Condensing: gas to liquid 8. Sublimation: solid to gas 9. Deposition: gas to solid 10. Enthalpy = heaet 11. 1 calorie = 4.18 Joules |
| Unit 9 |
Solution~ 1. Something dissolved in something else - homogeneous mizture of solute and solvent -Solute: what's being dissolved -Solvent: what's doing the dissolving Aqueous solution~solution where water is the solvent Suspension~ -Not a solution! (nothing is being dissolved) -Mixture of substances ion which on will settle to the bottom "Like Dissolves Like"~in terms of polarity -polar dissolves polar -nonpolar dissolves nonpolar -they do not dissolve each other Concentration~amount of solute per amount of solvent -to concetrate a solution add more solute -to dilute a solution, add more solvent Factors for Dissolving Rate~ 1. Particle Size - the smaller the particles the quicker they dissolve (more surface area) 2. Agitating the solution - stirring, swirling, shaking, etc. makes things dissolve faster 3. Temperature - an increase in temperature generally increases dissolving rate 4. Amount of solute already in solution - as a solution starts reaching its saturation point, the rate of dissolving begins to slow down (much like when you're eating and you are "starving" compared to beginning to feel full) Saturation~ 1. Saturated: solution holds maximum amount of solute that can be dissolved (no more solute can be dissolved at a specific temperature) 2. Unsaturated: solution doesn't hold the maximum (more solute can be dissolved at a specific temperature) -to calculate how much more can dissolve, subtract the saturation point from the amount of solute currently in solution Molarity~(M = moles of solute/liters of solution) -If grams are given, convert to moles using the molar mass before using the formula -if you need to convert to grams, solve for moles first then convert to grams using the molar mass Predicting Solubility~(whether or not the substance will dissolve in water) -If a substance is soluble, it is aqueous(aq). If a substance is insoluble, it is a solid(s), aka percipirate. -A precipitate is a solid that form from a reaction, and a precipitation reaction is a reaction that forms a precipitate. -Use solubility rules in reference table, pg 6 Solubility Curve~shows saturation points for different substances for a range of different temperatures -Grams of solute per 100g/100ml of water -Lines represent saturated solutions -Below the curve - unsaturated at that temperature -Above the curve - supersaturated at that temperature Equations~ 1. Molecular - balanced equation that shows ionic compounds as neutral instead of as ions (may be a precipitation reaction) 2. Ionic - shows how compounds becom ions as they dissolve in water -split up only AQUEOUS compounds into their two ions -leave solids in their compounds -spectator ions: ions that do not undergo change (ions as both reactants and products in the ionic equation) 3. Net ionic - shows the two ions that undergo change (to from a solid) |
| Unit 10 |
Equilibrium~rate of forward reactions = rate of reverse (speed at which reactants make products = speed at which products reform reactants) -Appears as though reaction stopped -Concentrations don't change (not necessarily equal; their ratios just don't change) Equilibrium Expression~ -[products]/[reactants]...raise to the power of their coefficients Le Chatelier's Principle~A system under stress will shift in a direction that relieves the stress Factors causing stress: 1. Pressure - increase in presuure shifts equilibrium to the side with fewer moles 2. Concentration - increase in reactants shifts towards products and vice versa 3. Absorption/release of heat -Endothermic: absorption of heat (heat is reactant) shifts to products -Exothermic: releas of heat (heat is product shifts to reactants |
| Unit 11 |
Nculear Decay~when a radioactive atom changes into a different isotope Alpha~ -Least penetrating -Release of a helium nucleus -Atomic mass decreases by 4 -Atomic number decreases by 2 Beta~ -More penetrating than alpha -Release of an electron (neutron mutates into a proton and electron - proton trapped in nucleus, electron emitted) -Atomic mass stays same -Atomic number increase by 1 (extra proton is trapped) Gamma~ -Most penentrating -Release of a wave of energy -Accompanies alpha or beta -No mass and no charge -Atomic mass and atomic number stay the same Half-Life~time required for half of a radioactive substance to decay Fission~splitting of atom -Happens in atomic bombs (uncontrolled) and nuclear reactors (controlled) -Produces a tremendous amount of energy -Chain reaction: products of one fission reaction cause other fission reactions to take place Fusion~combining two smaller nuclei to create a larger nuclei -occurs on the sun and in stars -produces a tremendous amount of energy Acid~produces H+ (or H3O+, hydronium) ions in water 1. Acids begin with H as the cation 2. Properties: -taste sour -turs litmus paper red -react with metals to produce H2 Base~produces OH- ions in water 1. Base ends with OH as the anion 2. Properties: -taste bitter -turns litmus paper blue -Feels slippery when wet Neutralization~Reaction of an acid and a base to produce a salt and water -Salt:compound formed when the cation of a base reacts with the anion of an acid Binary acid~composed of only 2 elements Monoprotic~one hydrogen Diprotic~two hydrogen's Triprotic~three hydrogen's Oxyacid~hydrogen, oxygen, and one other element Amphoteric~act as either an acid or base (such as water) pH~meausres acidity or basicity 1. Scale goes from 0-14 - >7 basic - <7 acidic - ~7 neutral 2. Formulas (on reference table, page 3) - pH= -log[H+] - pOH= -log[OH-] - pH + pOH = 14 - [H+]= 10-pH - [OH-]= 10-pOH Concentration vs. Strength of Acids/Bases~ -Concentration: how much acid/base is present in a given amount of solution -Strength: how completely the acid/base is ionized in water (how well they produce H+ or OH-) *Acids and bases are considered electrolytes because they break down to produce ions (H+ or OH-) in water. Strong acids and bases produce more of their respective ions in water (they are better conducters and ionize completely in water) while weaker acids and bases ionize partially to produce fewer of their ions. Naming Acids~ -No oxygen:add prefix hydro- and suffix -ic -Oxyacid: No prefix hydro-. If anion ends with -ate, change to -ic. If anion ends with -ite, change to -ous. Titration~addition of a known amount of solution (acids or bases) of KNOWN concentration to a known amount of solution (acid or base) of UNKOWN concentration -Trying to determine the concetration of either an acid or a base when you know the concentration of the other - M1V1 = M2V2...M = molarity; V = volume |
