EXAMBUSTERS MCAT CHEMISTRY

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Metric Abbreviations: Mega, Kilo, Hecto, Deka, Deci, Centi, Milli, Nano: Mega=M; Kilo=k; Hecto=h; Deka =da; Deci =d; Centi =c; Milli=m; Nano=n
Conversion Factor: A fraction which expresses an equality between two units of measurement and can be used to convert from one to the other (ex. 1kg/1000g)
Use conversion factors to solve: 1) How many kilograms in 2000 g? 2) How many feet in 60"?: 1) (2000g)(1kg/1000g) = 2kg; 2) (60")(1'/12") = 5 ft
Significant digits: Digits which are measured. All non-zero digits are significant. Zeros are significant unless they are placeholders
Scientific notation: A number expressed as Ax10B. "A" is between 1.00 and 9.99 and "B" is an integer.
Write conversion forumlas: 1) Celsius to Fahrenheit; 2) Celsius to Kelvin: 1) Ëšf = (1.8 x Ëšcelsius) + 32; 2) Kelvin = Ëšcelsius +273
Energy: The ability to do work; it is released or absorbed during chemical reactions in the form of heat, light, electricity. (calorie, Joule: 1cal = 4.18J)
Matter: A substance that occupies space has mass.
Compare Weight vs. Mass: Mass is the amount of matter. Weight measures gravitational force. Mass never varies. Weight can vary.
Compare potential energy and kinetic energy: Potential energy is energy due to position. KE is energy of motion.
Explain how to convert a number greater than 1 to scientific notation: Move decimal point to left until only 1 digit remains to left. Indicate number of moves as a positive exponent of 10. 3301 = 3.301x10^3
Explain how to convert a number less than 1 to scientific notation: Move decimal point to right until only 1 digit remains to left. Indicate number of moves as a negative exponent of 10. (0.00356 = 3.56 x 10-3
Explain the rules for multiplying the numbers in scientific notation: Multiply the first numbers and add the exponents. (3x10^5)(2x10^3)= 6x10^8
Explain the rules for dividing the numbers in scientific notation: Divide the first numbers and subtract the exponents. (8x10^6)/(2x10^10) = 4x10^-4
Physical properties: Properties which can be observed without changing the substance into something different. Color, odor, hardness, density, luster, state, conductivity, solubility, boiling and melting points
Chemical properties: A chemical property is observed when a substance changes into a new substance. Iron forms rust in air & water; gasoline burns in oxygen
Physical change: Substance changes form or state only. Boiling, melting, freezing, dissolving, grinding, cutting
Chemical change: Where new substances are formed with new chemical and physical properties. Oxygen & hydrogen form water; sodium & chlorine form sale (sodium chloride)
Density: Write the general equation and three standard units: The mass of a unit volume of a substance. Density = mass/volume = g/mL; g/L; kg/L
Element. List some examples: A substance is composed of identical atoms. Gold, silver, oxygen, hydrogen, lead, chlorine, helium, iron, copper, fluorine, arsenic
Compound. List some examples: substance composed of two or more elements chemically combined. Water - H2O; Salt - NaCl; Sugar - C6H12O6; Ammonia - NH3
Mixture. List some examples.: A combination of substances held together by physical means (dirt, milk, soup, saltwater, granite)
Homogeneous and Heterogeneous mixtures. Provide examples.: Homogeneous mixtures are uniform in composition (air, metal alloy, salt water). Heterogeneous mixtures are not uniform in composition (dirt, spaghetti sauce)
Three postulates of Dalton's Atomic Theory: 1) An element is composed of identical atoms; 2) Atoms of different elements have different properties; 3) Compounds are atoms of 2 or more elements chemically combined
The Law of Conservation of Mass: During a chemical reaction, matter is neither created nor destroyed
The Law of Constant Composition: A compound always contains the same elements combined in the same proportions by mass (H2O) is 88% oxygen no matter where it is found)
Law of Multiple Proportions: The same elements may combine to form more than one compound. The ratios of atoms are in small whole numbers (H2O and H2O2)
Atomic Mass Unit: the mass of a proton or neutron is equal to 1 atomic mass unit. Symbol - "amu"; 1 amu=1.66x10^-24
Atomic Number: What are the atomic numbers of helium, hydrogen, carbon, oxygen?: The number of protons in the nucleus of an atom of an element. Helium-2; carbon-6; hydrogen-1; oxygen-8
Mass number: The sum of protons plus neutrons n the nucleus of an atom
Isotope: Atoms which contain the same numbers of protons but different numbers of neutrons (ex. Hydrogen has 3 isotopes with mass numbers of 1,2,3)
Molecule: A group of two or more atoms held together by chemical bonds
Ion (provide examples): An atom or group of atoms which contains a positive or negative electrical charge (ex. Na+; Cl-; SO4^2-)
Cation and Anion (provide examples): cation - positively charged ion (Na+; Fe+2; NH4+; Ag+); anion - negatively charged ion (Cl-; SO4-2; OH-; P-3)
Valence electroncs: The electrons found in the outermost energy level of an atom
Oxidation number: A number (positive or negative) representing the charge on an ion or atom involved in a chemical bond
Three general rules for determining oxidation numbers: 1) Atoms of uncombined elements equal 0; 2) Hydrogen = +1 (in metallic hydrides =-1); 3) Oxygen = -2 (in peroxides =-1); (bonded with fluorine =+2)
Empirical Formula. What are the molecular and empirical formulas of peroxide?: An expression which gives the relative numbers of atoms of the elements in a molecule. Expressed as the lowest possible set of integers (H2O2, HO)
Molecular Formula. What is the molecular formula for ammonia?: An expression stating the number and kind of each atom present in a molecule of a substance (NH3 has 1 nitrogen atome and 3 hydrogen atoms in each molecule)
Positive Ion: which elements tend to form them?: Metals tend to form positive ions by losing electrons (Na â†’ Na+ e-)
Negative Ion: which elements tend to form them?: Non-metals form negative ions by gaining electrons (Cl +e- â†’Cl-)
Write oxidation numbers for ions of: Group IA & IIA; Group VIA & VIIA: IAâ†’+1; IIAâ†’+2; VIAâ†’-2; VIIAâ†’-1
Write formula for compounds of: 1) Sodium & Sulfate; 2) Magnesium & Nitrite; 3) Aluminum & Phosphate: 1) Na2SO4; 2) Mg(NO2)2; 3) AlPO4
Names of the ionic compounds: 1) FeCL3; 2) FeO; 3) Cu(OH); 3) Cu3PO4: 1) iron (III) chloride; 2) iron (II) oxide; 3) copper (II) hydroxide; 4) copper (I) phosphate
10 prefixes used to name covalent compounds: Mono=1; Di=2; Tri=3; Tetra=4; Penta=5; Hexa=6; Hepta=7; Octa=8; Nona=9; Deca=10
Forumlas and names of acids formed from: 1) F; 2) Cl; 3) Br; 4) I: 1) HF - hydrofluoric acid; 2) HCl - hydrochloric acid; 3) HBr - hydrobromic acid; 4) HI - hydriodic acid
Names of acids: 1)H2SO4; 2) HNO2; 3) H3PO4; 4)HClO: 1) sulfuric acid; 2) nitrous acid; 3) phosphoric acid; 4) hypochlorous acid
Write the equation for Percent Composition. What is the percent composition of Ca in CA(OH)2?: %=(me/FW)(100%); %= percent composition of the element; me=mass of element in one formula unit; FW = formula weight; (40/74)(100%) = 54% Ca in Ca(OH)2
Atomic Mass: A weighted average mass of the atoms of an element (assuming the carbon-12 isotope is exactly 12) (ex. Atomic mass of C=35.45 is calculated from two isotopes. Cl-35 and Cl-36)
Write the equation to calculate the atomic mass of an element: Atomic mass of X = ((mx1)(%x1)/100%) + ((mx1)(%x1)/100%)) + etc.; mx1, mx2, mxN = atomic masses of each isotope of element; %x1, %x2, %xN = percent composition of each isotope
Molecular Mass: Found by adding all the atomic masses of an element. Ex. H2O. Molecular mass = 18; H2 = 2(1) O = 16
Avogadro's Number: The number of atoms or formula units in "x" grams of an element or molecule where "x" is the atomic or molecular mass. (Always equal to 6.02 x 10^23)
Mole: 6.02x10^23 items, can be anything. The number of atoms in one mole (atomic mass in grams) of a monoatomic element. The number of formula units in one mole (formula mass in grams) of an ionic compound. The number of molecules in one mole (formula mass in grams) of a molecular substance)
For any substance, write a general formula to convert from Moles to Grams: g=(n)(MM) where, n = moles MM = molecular mass g = grams
For any substance, write a general formula to convert from Grams to Moles: n = g/MM; n = moles; MM = molecular mass; g = grams
For any substance, write a general formula to convert from Moles to Number of Particles: (n)(6.02x10^23) = P; n = number of moles; P = number of particles
For any substance, write a general formula to convert from Number of Particles to Moles: n = P/6.02x1023; n = number of moles; P = number of particles
Reactants; Products: 1) The starting materials in a chemical reaction; 2) The substances formed in a chemical reaction
Balance Equations: H2 + N2 â†” NH3; NaCL + Br2 â†” NaBr + Cl2: 1) 3H2 + N2 â†”2NH3; 2) 2NaCl + Br2 â†”2NaBr + Cl2
Diatomic; list 7 diatomic elements found in nature: A molecule composed of two atoms. H2, N2, O2, Cl2, F2, Br2, F2, I2
Synthesis or combination: A reaction where 2 or more elements form a compound; N2 + 3H2 â†”2NH3; 2H2 + O2 â†” 2H2O
Decomposition: A reaction where a compound breaks down into elements; CO2 â†” C + O2; 2CaO â†”2Ca + O2
Single Replacement: A reaction involving the replacement ina compound of an element by another element. Zn + CuCl2 â†” Cu + ZnCl2
Double replacement: A reaction where two ionic substances "trade" anions; NaCl + AgNO3 â†” NaNO3 + AgCl
Period: A horizontal row of elements in the periodic table. All have the same number of shells of e-. Across the period, the elements' properties change.
Column or Family: A verticle group of elements on the periodic table in the same column. They have similar properties and the same number of valence electrons.
Metalloid. List 5 examples.: An element with the properties of both metals and non-metals. Ex. Si, As, Ge, Sb, Te
Metals. List 6 characteristics: Solid (except Hg); malleable & ductile; conduct heat & electricity; shiny reflective & lustrous; lose e- to form cations; good reducing agents
Non-metals. List 7 characteristics: 1/2 are gases; solids are brittle; poor conductors of heat and electricity; dull & non-reflective; gain e- to form anions; good oxidizing agents
Transition element: A metal having two incomplete shells of electrons; many have multiple oxidation states; less active than family IA & IIA. Ex., Fe, Ag, Au, Cr, W
Describe the change in atomic radius across the periodic table: The atomic radius tends to decrease from left to right across the table and increase down the columns
Ionization energy: The energy change required for the removal of the outermost electron from the gaseous atom to form a +1 ion
Describe the change in Ionization Energy across the periodic table: Ionization energy tends to increase from left to right across the table and decrease down the columns
Electronegativity: The ability of an atom to attract to itself the e- in a covalent bond. Values range from 0.7 (Cs) to 4.0 (F)
Describe the change in electronegativity across the periodic table: Electronegativity tends to increase across the table left to right and decrease down the columns
Describe the change in metallic activity across the periodic table: Metallic activity decreases across the table (left to right) and increases down the columns
Describe the change in nonmetallic activity across the periodic table: Non-metallic activity increases across the table (L to R) and decreases down the columns
Compare the ionic radii of 1) Cl, S, P; 2) Na, Mg, Al; What is the reason for their different size?: 1) Cl-1 is smallest; P-3 is largest (the P atom gained the most e-); 2) Na+1 is largest Al+3 is smallest (the Al atom lost the most e-)
Quantum: A packet of energy associated with a specific wavelength of electromagnetic radiation
Quantum Number: A number used to describe the energy levels available to an electron. Each electron in an atom has a unique set of four.
Emission Spectrum: A bright line spectrum formed when energy absorbed by an element is emitted at specific wavelengths. Each element has a unique spectrum.
Absorption Spectrum: A dark-line spectrum formed when white light is passed through a vaporized element and a few specific wavelengths are absorbed.
List three basic postulates of the bohr model for the hydrogen atom: 1) e- are present only in specific energy states; 2) a quantum of energy is absorbed or emitted to change energy levels; 3) a quantum is the smallest amount of energy that can be gained or lost
Ground State: Electron is at its lowest energy level as close to the nucleus as possible
Excited State: An electron absorbs energy and moves to a higher energy level above the ground state
List four quantum numbers and their symbols: 1) Principal energy level "n"; 2) sublevel "l"; 3) orbital "m1"; 4) spin "ms"
List the names of the first four energy levels (or shells): 1 or K; 2 or L; 3 or M; 4 or N
List the names of the four sublevels and their electron capacities: sublevel s holds two electrons; p holds 6; d holds 10; f holds 14
Which sublevels are present in energy levels 1, 2, 3, and 4-7?: 1: s; 2: s, p; 3: s, p, d; 4-7: s, p, d, f
Describe the shapes of the orbitals: 1) s; 2) p; 3) d; 4) f: 1) s- sphere; 2) p - dumbbell shape with 2 lobes; 3) d - double dumbbell; most have 4 lobes; 4) f - most have 8 lobes
What is the order for filling sublevels (aufbau process) from lowest to highest energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d
Electron configuration: The distribution of electrons into shells and sublevels for an atom of an element. Each element has a unique electron configuration.
Write the electron configuration for 1) Lithium; 2) Iron: 1) 1s2, 2s1; 2) 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d6
Valence: The number of electrons in the atom's highest numbered shell.
What are the valences of the elements of families IA through VIIIA?: The valences of elements in groups IA - VIIIA are the element's column number. For example, the valence of Na is 1; O is 6
Electronegativity difference: A number found by taking the difference between the electronegativities of two atoms in a bond. Its value determines the type of bond.
Describe an ionic bond in terms of electronegativity difference: When the EN values differ by 1.7 or more. The atom with higher EN borrows the electrons from the atom with lower EN. The resulting positive and negative ions attract.
Nonpolar covalent bond: when the EN difference is very small (less than 0.5). Two bonded atoms share the valence e-. The resulting molecule has no electrostatic charge.
Polar covalent bond: When the EN difference is between 0.5 and 1.7, the bonding electrons stay closer to the more electronegative atom. Electrons are shared unequally.
Coordinate covalent bond: When both electrons in a covalent bond are supplied by one atom
Metallic Bonds: A sea of electrons surrounding positive metal ions
Hydrogen bonding: Formed when hydrogen is bonded to oxygen, fluorine, or nitrogen. The hydrogen of one molecule becomes attracted to the electronegative element of the other molecule. These intermolecular attractions cause higher boiling points than predicted
Octet Rule: Atoms tend to gain or lose outer shell electrons in order to achieve a noble gas configuration of 8 electrons
Double and Triple covalent bond: In a double bond, wo pairs of electrons are shared. In a triple bond, three pairs of electrons are shared.
Resonance structures: Where there is more than one possible bonding structure in a molecule
Hybrid orbitals (list three types): Where 2 or more pure atomic orbitals are mixed to form identical hybrid orbitals (ex. Sp, sp2, sp3)
Describe hybrid bonding in water, ammonia, methane: sp3 bonding results in a tetrahedron shape with bond angles of 109.5Ëš in methane and slightly less in water and ammonia.
Sigma bonds & Pi bonds: A sigma bond is present between any 2 orbitals except when 2 p orbitals share electrons; then this is a pi bond
List properties of ionic substances: Solids at 25C; Non-conducting as solids but conducting as aqueous solutions or liquids. Conducting as aqueous solutions or liquids. High melting & boiling points; Brittle; Low volatilities
List properties of molecular substances: Non-conducting as liquids and solids; Volatile liquids & solids; Many are gases at 25C; Low melting and boiling points; Soft and waxy solids
Exothermic: A chemical reaction which evolves heat
Endothermic: A chemical reaction which absorbs heat
Enthalpy: The heat content of a system
Enthalpy change (state the equation): The difference in heat content between the products and the reactants (Î”H = Î£Hproducts - Î£hreactants)
Exothermic reaction: Î”H is (positive or negative); Enthalphy is (increased or decreased): Negative; Decreased
Endothermic Reaction: Î”H is (positive or negative); Enthalphy is (increased or decreased): Positive; Increased
List in the order of increasing enthalpy: solid, gas, liquid: Solid (least enthalpy) -> Liquid -> Gas (most enthalpy)
Thermochemical equation: A balanced chemical equation which includes the enthalpy change. (H2(g) +1/2O2(g) â†’H2O(l) ; Î”H = -285kJ
In a thermochemical equation, what happens to Î”H when the moles of reactants double?: Enthalpy is directly proportional to mass. Therefore when the moles double, so does Î”H.
How are Î”H for a forward and Î”H for a reverse reaction related?: forward is equal in magnitude but opposite in sign to Î”H reverse.
Hess' Law: for a reaction is the same regardless of the path travelled from reactants to products.
Molar Heat of Formation: The molar heat of formation of a compound is equal to Î”H when 1 mold of compound is formed from its elements at 1 atm and 25C
Write the general equation to calculate Î”H for a chemical reaction: Î”H = Î£(Î”HF-products) - Î£(Î”HF-reactants); where Î”Hf = Heat of formation of reactants or products
Specific Heat. Give value for liquid water in calories and joules.: The amount of heat required to raise one gram of substance 1C. Water: 1cal/g-C or 4.18 J/g-C
Write an equation to calculate energy change when a fixed mass of substance changes temperature: Î”H = mCpÎ”T, where Î”T = temperature change, Cp = Specific Heat; Î”H = heat absorbed or given off
Charles Law: V1/T1 = V2/T2 (pressure and amount of gas are constant; V=volume; T=Kelvin)
Boyles Law: P1V1 = P2V2 (Temperature and amount of gas are constant, P=pressure, V=vol)
Combined Gas Law: P1V1/T1 = P2V2/T2 (amount of gas is contant. P=pressure, V=vol; T=Kelvin)
Dalton's Law of Partial Pressures: In a gas mixture, the total pressure equals the sum of the partial pressures of each component. Ptotal = P1 + P2 + P3⬦
Ideal Gas Law: PV = nRT, (P=pressure in atm; V = volume in L; n = # moles; R = 0.0820 Latm/Mol-K; T = Temp in K)
STP: Standard temperature (0C or 273 K) and Standard Pressure (1 atm or 760 torr)
Gay-Lussac's Law of Combining Gas Volumes: When only gases are involved in a reaction, the volumes of reactants and products are in a small, whole number ratio.
Avogadro's Law re. gases: Under the same conditions of temperature and pressure; equal volumes of gases contain equal numbers of moles
What is the volume of 1 mole of any gas at STP?: 22.4 L
Phase equilibrium: For a liquid in a closed container, when the rates of evaporation (liquid to gas) and condensation (gas to liquid) equalize; the concentration of each is stable.
Dynamic equilibrium: In a closed container where opposing changes are taking place at equal rates; the concentration of all components remains constant.
Normal Boiling Point: The temperature at which a liquid phase becomes a gas phase at a pressure of 1 atm.
Heat of Fusion (value for water): The amount of energy required to change a gram of substance from solid to liquid at its melting point (water = 80 cal/g)
Heat of Vaporization (value for water?): The amount of energy required per gram to change a liquid to a gas at its boiling point (water = 540 cal/g)
Triple Point: The only temperature and pressure combination at which the 3 phases of a substance (solid, liquid, gas) can co-exist in equilibrium
Vapor Pressure: The pressure the gas phase exerts on its liquid phase in a closed container. This pressure varies with temperature
Molarity - general equation: The number of moles of solute it a liter of solution; M = n/L
Solute (provide example): The substance dissolved in another (solvent). Salt is the solute in salt water.
Solvent (provide example: A substance, usually a liquid, into which another substance (solute) is dissolved. Water is the solven in iced tea.
Solution (provide example): A liquid, gas or solid phase containing 2 or more components uniformly dispersed (air, coffee, saltwater)
Solubility curves: A curve for a given substance which shows how much dissolves in a given amount of solvent at different temperatures.
How do temperatures and pressure affect the solubility of a solid?: Solubility usually increases with increasing temperature. Pressure has little effect.
How do temperature and pressure affect the solubility of a gas?: Solubility usually decreases with increasing temperature. Solubility increases in direct proportion to an increase in pressure.
Which three factors affect the rate of solubility?: Pulverizing; stirring; heating
What is a general rule for solubilities of polar and nonpolar compounds?: "Like dissolves like"; Ionic and polar solvents dissolve ionic, polar solutes (water dissolves salt). Non polar solvents dissolve nonpolar solutes (acetone dissolves gasoline)
List some basic facts about solutions: Particle size less than 1 mmicron; Clear (may be colored); Particles don't settle; Can pass through membranes; Particles not visible
List some basic facts about colloids: Particles measure 1-100 mmicrons; Particles don't pass through a membrane; Show brownian motion and the Tyndall effect; Particles don't settle; Clear and pass through filter paper
List some basic facts about suspensions: No brownian motion; Don't pass through filter paper or a membrane; Cloudy but particles settle on standing; Particles visible with microscope or eye
How many grams of NaCl are required to prepare 500 grams of a 5% solution?: %Concentration = (gNaCl/gsolution)(100%); 5% =(x/500g)(100%);x=25 g NaCl
Calculate the molality of 10 moles of H2SO4 dissolved in a 4 kg of water: Molality = Moles Solute / kg solvent = 10/4 = 2/5 Molal
Gram-equivalent weight: The amount of substance which reacts with or displaces 1 mole of H+ ions.
Normality: The number of gram-equivalent weights in a liter of solution
in H2O solutions: 1) How many ËšC is the freezing point depressed for each molal of solute? 2) How many ËšC is the boiling point elevated for each molal of solute?: 1) 1.86C for each molal of particles of solute; 2) 0.51C for each molal of particles of solute
List five factors that control reaction rate: nature of reactants; exposed surface area; concentrations; temperatures; presence of catalyst
State the collision theory of reaction rates: There must be collisions between reactants. Reaction rate depends on number of collisions per unit time and the percent which are successful (Have sufficient energy)
How is the reaction rate related to concentration?: Reaction rate is directly proportional to the concentrations of reactants
Activation energy: The energy necessary for a reaction to begin. Obtained from the kinetic energy released during collision
Catalyst: a catalyst is introduced into a reaction to speed it up or slow it down. It is not consumed. An increase or decrease of activation energy results from an alternate reaction path.
Law of Mass Action: The rate of a chemical reaction is proportional to the product of the concentrations of the reactants
Specific Rate Constant: Symbol is "k" in a rate equation. A constant specific to temperature and reaction which is part of every rate equation
Reversible reaction: A system where the following opposite reactions are taking place: reactant becoming product; product becoming reactant
Equilibrium: The point in a reversible reaction where the forward and reverse reactions are taking place at the same rate.
Are concentrations of product and reactant equal at equilibrium?: No. The are constant but not equal. Their relative concentrations are determined by the value of the equilibrium constant at that temperature.
Write the equilibrium expression for aA + bB â†” cC + dD: Keq = [C]^c[D]^d / [A]^a[B]^b, where Keq = Equilibrium constant
How are reactant & product concentrations related to the magnitude fo Keq?: Keq is large: [reactant] is small and [product] is large; when Keq is small: [reactant] is large and [product] is small
Le Chatelier's Principle: If stress is placed on a system at equilibrium, the equilibrium shifts in order the counteract the effects of the stress and regain equilibrium
How does a concentration change affect equilibrium?: If one of the substances is added or removed, all the concentrations of substances adjust to a new equilibrium with the same Keq
How does a change in temperature affect equilibrium: The reaction shifts to a new equilibrium point with a new Keq. If the temperature is raised, the equilibrium is shifted to reaction which absorbs heat.
How does a pressure change affect equilibrium?: Only in reactions where gases are involved. The reaction will shift to oppose pressure change, resulting in fewer moles of gas particles
Ionization constant. Write the expression for the ionization of acid "HA": For substances in solution that partially ionize. An equilibrium expression may be written with Ki; Ki = [H+][A-] / [HA]
What are the two driving forces that control reactions?: A drive towards increased entropy (disorder). A drive towards decreased enthalpy (lower heat content)
Second Law of Thermodynamics: The entropy of the universe increases for any spontaneous process.
Free Energy Change (Î”G). Write the free energy equation.: A property which reflects a system's capacity to do useful work. Î”G = Î”H - TÎ”S; G=free energy; S=entropy; H=enthalpy; T=kelvin
How can Î”G be used to predict if a reaction is spontaneous?: When Î”G is positive it is not spontaneous; negative, it is spontaneous; equals 0 it is at equilibrium.
Ionization Constant: Kw = 1 x 10-14 at 25C
pH. For what values is a solution acidic, basic and neutral?: pH = -log[H+] ; The degree of acidity of a solution. <7 =acid; 7 = neutral; >7 = basic.
pOH. For what values is a solution acidic, basic, neutral?: pOH = -log[OH-]; The degree of basicity of a solution. <7 = basic; 7 = neutral; >7=acid.
How are pH and pOH of a solution related?: The sum of the two values = 14; (pH +pOH = 14)
How are the concentrations of [H+] and [OH-] related in a solution?: The product of the concentrations = 1x10-14; [H+][OH-] = 1x10-14
Solubility Product Constant (Ksp): An equilibrium exists in a saturated solution between dissolved and undissolved solute. Ksp is the equilibrium constant for this reaction.
Write the solubility product expression for AgCl â†” Ag+ + Cl-: Ksp = [Ag+][Cl-]
Six common characteristics of acids: Form H2O solutions; Conduct electricity; React with active metals; Turn blue litmus red; Neutralize bases; Sour taste
Seven common characteristics of bases: Form H2O solutions; Conducts electricity; Turns red litmus blue; Feels slippery; Caustic; Neutralizes acids; Bases + fats form SOAP
Neutralization Reaction - Write equation for hydrochloric acid and sodium hydroxide: acid + base â†’ salt + water; HCl + NaOH â†’NaCl +H2O
Arhennius Theory: An acid yields protons in solution (H+ ions); A base yields hydroxide ions in solution (OH- ions)
Bronsted Theory: An acid is a proton donor. A base is a proton acceptor.
Lewis Theory: An acid is an electron pair acceptor; a base is an electron pair donor.
Conjugate Base - Write conjugate base of HCl: When a Bronsted acid donates a proton, it becomes its conjugate base (conjugate base of HCl is Cl-)
Conjugate Acid - Write conjugate acid of I-: When a Bronsted base accepts a proton it becomes its conjugate acid (conjugate acid of I- is HI)
Calculate the volume of 10M NaOH needed to titrate 5L of 2M HCl: MA x VA = MB x VB; 2M x 5L = 10M x VB; 1L = volume of base
Electrolyte (give 3 examples): A solute whose aqueous solution contains ions and conducts electricity (acids, bases, salts)
Nonelectrolyte (give 2 examples): A solute whose aqueous solution does not conduct electricity (sugar, benzene, most organic compounds)
Anode: A positively charged electrode which attracts anions. Where oxidation takes place.
Cathode: A negatively charged electrode which attracts cations. Where reduction takes place.
Oxidation: The loss of electrons (ex. Cumetal â†’Cu+2 +2e-
Reduction: The gain of electrons (Zn+2 +2e- â†’ Znmetal)
Electrode Potential: A measure in volts of the tendency of atoms to gain or lose electrons. (Relative to a H2 oxidation reaction which has an assigned value of zero)
Half Reaction: One of the two parts, either the reduction or the oxidation, of an oxidation-reduction reaction
How can you determine if a redox reaction will take place spontaneously: Add the electrode potentials of the two half reactions. If the result is positive, the reaction is spontaneous; if negative, the reaction is not spontaneous.
Faraday: A unit of electric charge which deposits by electrolysis one equivalent weight of an element. Equals 96,500 coulombs
Equivalent weight: The number of grams of an element which will accept or donate 1 mole of electrons
Oxidizing agent: A substance which causes another substance to be oxidized (oxidizing agent is simultaneously reduced)
Reducing agent: A substance which causes another substance to be reduced (reducing agent is simultaneously oxidized)
List some basic facts about carbon bonding: Carbon forms more compounds than any other element; Each atom requires 4 covalent bonds; Carbon can form long chains and rings; Bonds commonly to O, H, N, S, P, and halogens
Alkane: A series of hydrocarbons with only single covalent bonds (CnH2n+2)
Alkene: A series of hydrocarbons containing at least one double covalent bond (CnH2n)
Alkyne: A series of hydrocarbons containing at least one triple covalent bond (CnH2n-2)
List prefixes for naming hydrocarbons for 1-10 carbons in a molecule: 1 meth; 2 eth; 3 prop; 4 but; 5 pent; 6 hex; 7 hept; 8 oct; 9 non; 10 dec
List names for alkanes with 1 to 10 carbons in a molecular chain: 1 methane; 2 ethane; 3 propane; 4 butane; 5 pentane; 6 hexane; 7 heptane; 8 octane; 9 nonane; 10 decane
Substitution reaction. Which elements commonly substitute in alkanes?: A reaction where one of the hydrogen atoms in a hydrocarbon is replaced by another. Usually a halogen. (CH4 + Br2â†’Ch3Br + HBr
Cycloalkane: An alkane which has a ring structure instead of a chain
Saturated and Unsaturated. Which are sturated? Alkanes, alkenes, alkynes.: A compound only containing single covalent bonds is saturated. Alkanes are saturated. A compound containing double or triple bonds is unsaturated. Alkenes and alkynes are unsaturated.
Addition reaction: In an unsaturated hydrocarbon, two atoms may be added to the structure across a double or triple bond (C2H2 + Br2 â†’ Ch2Br2)
Aromatic compounds: Unsaturated ring structures with six carbon atoms. Benzene is the simplest aromatic. (CnH2n-6)
Isomer: Compounds with the same molecular formula but different structural formulas (different connectivity)
Polymerization: The combination of two or more unsaturated molecules to form a larger chain molecule. This is how plastics are made.
Hydrogenation: The process of adding hydrogen to an unsaturated hydrocarbon
Dehydrogenation: The process of removing hydrogen from a hydrocarbon
Alcohol: hydrocarbons which contain the hydroxyl functional group (OH-) attached to a saturated carbon (R-O-H)
Aldehyde: A hydrocarbon containing the aldehyde functional group (R-C(=O)-H)
Carboxylic Acid: A hydrocarbon containing the carboxyl functional group. (R-C(=O)-O-H)
Ketone: A hydrocarbon containing a ketone functional group (R-C(=O)-R')
Ether: A hydrocarbon containing an ether functional group. (R-O-R')
Amine: A hydrocarbon containing an amine functional group. (R-NH2)
Ester: A hydrocarbon containing an ester functional group. (R-O-C(=O)-R')
Amino Acid: Organic compounds which contain an amine and a carboxyl group. (H2N-CH(-R)-COOH)
Compare primary and secondary alcohols: Primary: The OH- group is attaced to the end carbon of the chain. Secondary: The carbon bearing theOH- goup is directly attached to two other carbons.
Alpha particle: a helium nucleus; charge =+2; High energy; Low velocity; Ejection reduces atomic number by 2amu and atomic weight by 4amu
Beta particle: An electron ejected from the nucleus when a neutron decays to a proton; Increases atomic number by one; High velocity; Low energy
Gamma radiation: Usually emitted with beta radiation; Has neither charge nor mass; High energy; Travels at the speed of light
Half-life: The time it takes for half of a radioactive sample to decay. It can range from a fraction of a second to many years.
Transmutation: The conversion of an element to a new element due to a change in number of protons. Ex. Alpha or beta decay
Fission: The breakdown of heavy nuclei into lighter nuclei. The source of nuclear power
Fusion: The joining of lighter nuclei to form heavier nuclei. Source of the sun's energy
Write the nuclear equation: 1) U238 loses an alpha particle; 2) Th234 loses a beta particle.: 92U238 â†’90Th234 +2He4; 90Th234â†’ 91Pa234 + -1e0

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