Final Exam Study Guide

Chemistry Final Exam Study Guide Acids and Bases Acid: substance that provides H+ ions in water Base: substance that provides OH- ions in water              Covalent Bonds Review: ionic bonds are formed by the positive charge of the cation being attracted to the negative charge of the anion Covalent bonds: bonds formed by sharing electrons between the atoms (not ions) o 1 electron from each atom (2 atoms per bond) o Electrons available for bonding are the un-paired electrons in orbitals (orbitals with only 1 electron in them) Molecule o A group of atoms held together by covalent bonds Atoms will bond with enough other atoms to complete the octet o But they can only form as many bonds as unpaired valence electrons Hydrogen has 1 valence electron, so it can only form 1 bond, and will not achieve an octet Boron has 3 valence electrons, so can only form 3 bonds, and will not achieve the octet (only 6 electrons) o There are always exceptions to the rules Sulfur and phosphorus can extend the octet Multiple Covalent Bonds between 2 atoms If the octet is not met, additional bonds may be formed o Co2 (carbon dioxide) C has 4 valence electrons, so 4 bonds possible O has 6 valence electrons, 2 are unpaired, so 2 bonds are possible PRACTICE IONS: Sodium o Symbol: Na o Valence: 1 o Electron dot: Na* o Ionic dot: Na+ (loses its electron to get a 1+ charge) Magnesium o Symbol: Mg o 2 valence Eo Loses 2 electrons to get 2+ charge Oxygen o Symbol: O o Valence: 6 o Will gain 2 electrons for a 2- charge Cannot have a covalent bond with ions, just atoms  Expanding Condensed Structures Connectivity is from left to right Drawing Lewis Structures 1. Determine the total number of valence electrons in the molecule or ion 2. Draw a line between each pair of atoms connected. Each line represents 2 electrons. 3. Use the remaining electrons on the outer atoms to satisfy the octet rule 4. Place any remaining electrons in lone pairs around the central atom 5. If the central atom does not satisfy the octet, use lone electrons from the outer atoms to form additional bonds between the center and outer atoms.   Resonance Structures If more than one valid Lewis structure can be drawn for a molecule or ion, then the molecule or ion has resonance, and each of the possible Lewis structures are resonance structures of each other Resonance structures are identified by a double-ended arrow pointing between each structure   Number of resonance structures Look at the number of equivalent atoms surrounding the center atom o If the bonds between each of the equivalent atoms and the center atom are the same, there are no resonance structures o If the bonds between each of the equivalent atoms and the center atom are NOT the same, there ARE resonance structures The number of resonance structures is determined by how many spots on the molecule you can move the different bonds to        VSEPR (Valence Shell Electron Pair Repulsion) Model used to predict the molecular shapes and geometries of molecules Identifies the atoms and lone pairs of electrons connected to an atom of interest as an "electron charge cloud" AXE designation: o AXaEb Where a=number of atoms connected to atom of interest Where b= number of lone pairs of electrons on the atom of interest A= center atom X= other atoms attached to center atom  E= lone pairs of electrons Geometries of molecules 1. Draw the most stable Lewis structure for the molecule 2. Count the number of electron charge clouds around the atom of interest. An electron charge cloud can be 1 lone pair of electrons or 1 connected atom 3. The geometry corresponds to the number of electron charge clouds around the atom of interest a. 2 electron charge clouds= linear b. 3 electron charge clouds= trigonal planar c. 4 electron charge clouds= tetrahedral d. 5 electron charge clouds= trigonal bipyramidal e. 6 electron charge clouds= octahedral Shapes of molecules 1. Draw the most stable Lewis structure for the molecule 2. Determine the geometry of the molecule 3. Identify how many electron charge clouds are lone pairs of electrons and how many are atoms a. If all are atoms, then geometry=shape b. If not, move to the next step 2. If lone pairs of electrons exist Geometry Number of Atoms Number of Electron Pairs Shape AXE AX2E0 Linear 2 0 Linear Trigonal planar 3 0 Trigonal Planar AX3E0 2 1 Bent AX2E1 4 0 Tetrahedral AX4E0 3 1 Trigonal pyramidal AX3E1 2 2 Bent AX2E2 Tetrahedral Lecture 12: Scientific notation: a compact way to write very large or very small numbers - Units o Specific units of measurement are defined by the International System of Units o Ex. Meters for length/distance, kilograms for mass, kelvin for temperature Units and Equivalents Quantity SI Unit (symbol) Mass Kilogram (kg) Length Volume Meter (m) Cubic meter (m3) Temperature Kelvin (K) Metric Unit (symbol) Gram (g) Meter (m) Liter (L) Celsius degree © Equivalents 1kg=1000g 1kg=2.205lb 1m=3.280ft 1 meter cubed=1000L 1 meter cubed=264.2 gal See section 1.11 Time Seconds (s) Seconds (s) - Metric Prefixes and Scientific Notation - Metric prefixes o Used to express very large or very small values  Ex. Easier to express 1,000,000 grams as 1 megagram (Mg)  Ex. Easier to express .000000001 meters as 1 nanometer (nm) or in scientific notation Common Metric Prefixes Prefix Symbol mega M Base Unit Multiplied By 1,000,000 kilo hecto Deka Deci Centi Milli Micro Nano Picto Femto Mass vs. Weight k h da d c m µ n p f 1000 100 10 0.1 0.01 0.001 0.000001 0.000000001 0.000000000001 0.000000000000001 - Example 1 megameter=10,000,000 m 1 kilogram=10,000 g 1 hectogram=100g 1 dekaliter=10L 1 deciliter 1 centimeter=0.01m 1 milligram=0.001 g 1 micrometer=10-6m - Mass is the measured amount of matter in an object Weight is the measured gravitational forces that the earth or other large bodies exert on an object When we commonly talk about how much we weight, or the weight of things, we are actually talking about what their mass is Mass Conversions Unit 1 kilogram (kg) 1 gram (g) 1 milligram (mg) Equivalent =1000 grams =2.205 pounds =0.001 kilogram =1000 milligrams =0.03527 ounces =0.001 gram =1000 micrograms Unit 1 ton 1 pound (lb) 1 ounce Equivalent =2000 pounds =907.03 kilograms =16 ounces =0.454 kilogram =454 grams =0.02835 kilograms =28.35 grams =28,350 milligrams =0.000001 gram =0.001 milligram Common units of length - The US commonly uses inches, feet, and miles whereas other countries and their scientific community use millimeters, centimeters, meters, and kilometers (See table 1.8 in book for units and conversion factors) Units of Volume Unit 1 cubic meter (m3) 1 liter (L) 1 deciliter (dL) 1 milliliter (mL) 1 microliter (µL) 1 gallon (gal) 1 quart (qt) 1 fluid ounce (fl oz) Equivalent =1000 liters =264.2 gallons =0.001 cubic meters =1000 milliliters =1.057 quarts =0.1 liter =100 milliliters =0.001 liter =1000 microliters =0.001 milliliter =3.7854 liters =0.9464 liter =946.4 milliliters =29.57 milliliters Lecture 13 Significant Figures - Significant figures and scientific notation are used to represent the precision of a measurement When in doubt, rewriting the number in scientific notation can clarify which of the figures are significant Rules: 1) Zeroes in the middle of the number are always significant (ex. 1.0023 5 sig figs) 2) Zeroes at the beginning of a number are not significant (ex. 0.00123 3 sig figs) 3) Zeroes at the end of a number AFTER the decimal point are always significant (ex. 1.23000 6 sig figs) 4) Zeroes at the end of a number BEFORE the decimal point or implied decimal point may or may not be significant (ex. 1200 without a decimal point 2 sig figs)     1 microgram (µg) Determining Significant Figures for Calculation Results  - - Rounding your answers to the correct number of significant figures  -    Multiplication/Division o The number of significant figures in your answer is equal to the smallest number of significant figures in your original numbers Conversion factors do not count, they are considered exact numbers Addition/Subtraction o Then number of significant figures in your answer is determined by the right-most decimal place that all of your original numbers carry Rule 1: If the first digit that you remove is 4 or less, do not round your last kept digit o Ex. Round 1.4539 to 3 sig figs 1.45 Rule 2: If the first digit that you remove is 5 or greater, round your last kept digit up by 1 o Ex. Round 6.1497 to 3 sig figs 6.15 Temperature, Heat, Energy, and Density - - Temperature: the measured amount of heat energy in an object o Units (K, degrees F, degrees C) Energy: the capacity to do work or supplyheat o Units (cal, kcal, J) Specific heat: the amount of heat required to raise the temperature of 1g of substance by 1 degrees C (AKA heat capacity) o Q=mxCxΔT where q=heat, m=mass, c=specific heat and ΔT=Tfinal-Tinitial Density: the physical property relating the mass of a per specific volume of a substance o Unit: (g/mL) Lecture 14 Chapter 5: Chemical Reactions - Reactants: starting materials/substances in a reaction. Located on the left-hand side of a reaction equation Products: materuals/substances formed from the reaction of starting materials. Located on the right-hand side of a reaction equation Reactants products Law of conservation of mass: matter is neither created nor destroyed, so elements must balance on both sides of the reaction equation Balancing Reaction Equations - Stoichiometric Coefficient: number placed in front of a chemical formula in order to balance your reactants and products o A+BA2B3 (not balanced) o 2A+3B A2B3 (balanced) Steps for Balancing 1. Write the correct chemical formulas for all reactants and products, placing reactants on the left, products on the right. 2. Add appropriate coefficients to balance the number of atoms of each element 3. Check to make sure the number of each element matches on both sides of the reaction 4. Make sure your coefficients are reduced to the lowest whole number values. Precipitation Reactions - Precipitate: insoluble solid that forms in solution. Phase is (s). Solubility: the amount of compound that will dissolve in a given amount of solvent at a given temperature. Solubility rules: indicate which combinations of ions are insoluble (form precipitates) or soluble (do not form precipitates). Aqueous solutions: solutions made from dissolving soluble compounds in water. Phase is (aq). Lecture 15 Acid and Base Reactions - Neutralization reactions: an acid and a base react to form an aqueous salt and water Anion from the acid and cation from the base form the salt. Redox Reactions - - Redox=reduction-oxidation Redox reaction: reactions in which electrons are transferred from one atom to another. Reduction: the gain of one or more electrons to an atom Oxidation: the loss of one or more electrons from an atom o Think of electrons as products or reactants in determining if the atom is gaining or losing them. Classified as any reaction where an atom is in its elemental state on one side of the reaction and is in a compound or ionic form on the other. Oxidation Numbers/Guidelines - Elemental states have oxidation numbers of zero (so all single elements are 0) A monoatomic ion has an oxidation number equal to its charge In a molecular compound, the atom typically has the same oxidation number as it would as a monoatomic ion, but charges must add up to 0. Ionic and Net Ionic Equations - Ionic equations show all aqueous ionic compounds broken up into their ions Net ionic equations remove any spectator ions o Ions that appear in the same form on both sides of the reaction equation Lecture 16 Chapter 6: Mole and Mass - Molecular mass: sum of atomic masses of all atoms in a molecule (unit=amu) Formula mass: sum of atomic masses of all atoms in one formula unit of any compound, whether ionic or molecular Molar mass: mass in grams of 1 mole of a substance. Same number as molecular mass but units are different (grams instead of amu) Avagadro’s number: the number of formula units in 1 mole of ANYTHING o 6.022x1023 Lecture 23 Term Spontaneous Non-spontaneous Spontaneity and change of Enthalpy (ΔH) Entropy (s) Change in entropy (ΔS) Spontaneity and Entropy (ΔS) Gibbs Free Energy (ΔG) Definition Occuring without external influence Must have a continuous external influence in order to occur Exothermic favors spontaneous Endothermic favors non-spontaneous The measure of chaos (disorder) in a system - Gases have high entropy, solids have low entropy The change in entropy for a reaction - A change from a low entropy state to a high entropy state has a positive ΔS and high entropy to low entropy has a negative ΔS A positive ΔS favors spontaneity Combining all the effects of ΔH and ΔS to determine the overall spontaneity of a reaction by examining Gibbs Free Energy ΔG= ΔH-T ΔS where G= gibbs free energy, H= heat of reaction, T= temperature in kelvin, and S= entropy change - - A negative ΔG means a is a spontaneous reaction (exogenic) A positive ΔG means a non-spontaneous reaction (endogenic) The value of ΔG for a reaction changes sign when the reaction is reversed (same value, opposite sign) Some reactions can have a change in spontaneity with a change in temperature Lecture 25 Le Chatlier’s Principle - When a stress is applied to a system at equilibrium, the equilibrium shifts to relieve that stress Concentration Temperature Pressure Catalysis If you increase the concentration of a reactant, equilibrium shifts towards products. If you increase the concentration of a product, equilibrium shifts towards reactants. If endothermic, heat is a reactant, so increase temperature increases a reactant, and shifts equilibrium towards reactants If exothermic, heat is a product, so increasing the temperature increases a product, and shifts equilibrium towards reactants Increase in pressure shifts equilibrium towards the side with fewer moles of gas Decrease in pressure shifts equikibrium towards the side with more moles of gas Equilibrium is reached faster, but K does not change Lecture 26 Chapter 8: Gases, Liquids, and Soilods Term Phase change Melting Freezing Vaporization (boiling) Condensation Sublimination Deposition Definition Matter changing from one phase to another Phase change from solid to liquid Phase change from liquid to solid Phase change from liquid to gas Phase change from gas to liquid Phase change from solid to gas, bypassing liquid Phase change from gas to solid, bypassing liquid London Dispersion Weakest intermolecular force. ALL molecules experience London Dispersion (LD) forces Typically stronger than LD forces. ONLY occur between polar molecules The strongest intermolecular force. ONLY occurs when molecules have an N-H, O-H, or F-H bond Dipole-dipole Hydrogen bonding Kinetic Molecular Theory and Gases Kinetic Molecular Theory is used to describe the behavior of gases: - Gases consist of many particles moving around at random with no attractive forces Gas particles occupy much smaller space between particles The average Kinetic Energy (KE) of gas particles is proportional to the temperature in Kelvin All gas particle collisions are elastic. NO energy is lost during collisions, so KE particles remain constant. Gases that obet ALL of the conditions from the Kinetic Molecular Theory are considered IDEAL gases Most real gases display nearly ideal behavior Pressure - Defined as a force applied per unit area P=F/A Most common unit is atmospheres (atm), others include : o Millimeters Mercury (mmHg) 760 mmHg=1atm o Torr 760 torr=1atm o Pound per square inch (psi) 14.7psi=1atm o Pascal (Pa) 101325 Pa=1atm     Intermolecular Forces Lecture 27 Gas Laws Name Boyle’s Law (constant T) Equation P1V1=P2V2 Charle’s Law (constant P) V1 V2 = T 1 T2 Law Pressure is inversely proportional to volume at constant temperature Volume is directly proportional to temperature at a constant pressure Gay-Lussac’s Law (constant V) P1 P2 = T1 T 2 Combined Gas Law (constant amount) P1 V 1 P2 V 2 = T1 T2 Avagadro’s Law (constant T and P) V1 V2 = n 1 n2 Ideal gas law (constant R) PV=nRT Dalton’s Law of Partial Pressure Ptotal= Pa + Pb + Pc… %gas= P gas x 100% Ptotal Pressure is directly proportional to temperature at constant volume Relates pressure, volume, and temperature where the amount of gas stays constant Relates volume to the molar amount at a constant temperature and pressure Relates pressure, volume, temperature, and molar amount to the ideal gas constant R R= 0.08206 Lxatm/molxK The total pressure of a gas mixture equals the sum of the partial pressure of the gas mixture components Lecture 28 Liquids, Solids, and Phase Diagrams Term Vapor Vapor pressure Definition Where gas molecules are in equilibrium with a liquid The partial pressure of a vapor (gas) molecules in equilibrium with a liquid - Vapor pressure increases as temperature increases - The stronger the intermolecular forces, the lower the vapor pressure - The stronger the intermolecular forces, the higher the boiling point Properties of Liquids Term Viscosity Surface Tension Definition A liquid’s resistance to flow - High viscosity flows slowly (strong intermolecular forces) - Low viscosity flows quickly and freely (weak intermolecular forces) The resistance for a liquid to spread and increase its surface area - High surface tension liquids keep to a small area - Low surface tension liquids spread out as - much as possible Liquids want to limit the number of surface molecules because they are less stable and experience less intermolecular forces with other molecules in the liquid. Interior molecules experience more intermolecular forces and are more stable Crystalline Solids - Solids that have an ordered arrangement/orientation of particles over long range. Particles can be ions, molecules, or atoms. o Ionic solids: composed of ions  Brittle, high melting points o Molecular solids: composed of molecules  Soft, low to moderate melting points o Covalent solids: composed of covalently bonded molecules  Very hard, very high melting points o Metallic solids: composed of metal cations surrounded by a sea of electrons  Shiny (lustrous), varying from soft to hard, high melting poits Amorphous Solids - Solids whose orientation/ arrangement is random o Shiny (lustrous) o Vary from soft to hard o High melting point Phase change - Energy involved in changing the state (phase) of a substance o Heat of fusion: the quantity of heat required to completely melt 1 mole of a substance one it has reached its melting point o Heat of vaporization: the quantity of heat required to completely vaporize 1 mole of a substance once it has reached its boiling point  Lecture 24 Term Activation energy (Ea or Eact) Rate of Reaction Definition The energy required for a reaction to occur The measure of how quickly a reaction occurs. Low activation energy reactions occur quickly in comparison to high activation energy reactions. An increase in temperature increases frequency and forcefulness of collision, therefore increases the reaction rate An increase in concentration increases the frequency of collisions, therefore increases the reaction rate Adding a catalyst lowers activation energy, therefore increases the rate of a reaction Temperature Concentration Catalysis Reversible Reactions Idenfitifed by reversible arrows instead of or  A reaction read from left to right is the forward reaction Reading the same reaction from right to left is the reverse reaction The rates of the forward and reverse reactions do not have to be the same If the rates of the forward and reverse reactions ARE the same, then the reaction is at equilibrium  - Lecture 25 Equilibrium and Equilibrium Constant - K or Keq A value that describes the equilibrium state of a reaction for aA+bB cC+dD [C ]c [D ]d K= a b [ A ] [B] K much smaller than 0.001: Only reactants are present at equilibrium in any significant amounts, so essentially no reaction occurs K between 0.001 and 1: more reactants than products are present at equilibrium K between 1 and 1000: More products than reactants are present at equilibrium K much greater than 1000: Only products are present at equilibrium in any significant amounts, so reaction essentially goes to completion Lecture 26 Term Phase change Melting Freezing Definition Matter changing from one phase to another Phase change from liquid to solid Phase change from liquid to solid Vaporization Condensation Sublimation Deposition Phase change from liquid to gas (also known as boiling) Phase change from a gas to a liquid Phase change from a solid to a gas, bypassing liquid Phase change from a gas to a solid, bypassing liquid Lecture Term Acid Base Neutralization Hydronium ion Bronsted-Lowry Acid Bronsted-Lowry base Definition A substance that produces hydrogen ions (H+) when dissolved in water. Generally called HA A substance that produces hydroxide ions (OH-) when dissolved in water. Generally called B When an acid and a base react to form water and an aqueous salt H30+; formed when water reacts with an acid A substance that is able to donate (give off) a hydrogen ion A substance that is able to accept (take) a hydrogen ion Conjgates Removing the H+ ion from an acid results in the acid’s conjugate base Adding an H+ ion to a base results in the base’s conjugate acid Acid and base strength - The stronger the acid or base the more the acid or base dissociates (breaks up into ions) when dissolved in water Weak acids and bases do not fully dissociate into ions Strong acids and bases DO fully dissociate into ions The stronger the acid/base, the weaker the conjugate base/acid Lecture 35 Because water can act as both an acid and a base, we can use it’s dissociation into ions to calculate the relative amounts of OH- and H+ ions - The larger the H3o+, the more acidic a solution is - If H3o+ is large, then OH- is small pH scale - Converts H3o+ from exponential notation to a scale of 0-14 using the negative log pH= -log[H3o+] or pH= -log{H+] Acidic solitions: pH<7 Neutral solutions: pH=7 Basic solutions: pH.7 Normality of Acid and Base - Uses the molarity with the molecular formula to determine the concentration of [H+] or [OH-] Normality of acid = [acid]x # H+ ions per formula unit Normality of base= [base]x # OH- ions per formula unit Normality gives you the [H+] or [OH-] Lecture 36 Term Neutralizations Co2 (g) producing reactions Ammonium salt producing reactions Acidity and Basicity of Salt Solutions Definition Reactions with hydroxides Reactions with hydrogen carbonates or carbonates Reactions with nitrogen-containing compounds (predicting the acidity/basicity of a salt solution based on the cation and anion of the salt) Anion Derived from Acid that is: Strong Weak Strong Weak Cation derived from base that is: Weak Strong Strong Weak Solution Acidic Basic Neutral More information needed Buffer Solutions - Buffers are conjugate acid base pairs that are resistant to change in pH o Typically a weak acid and its conjugate base o The conjugate base can be paired with metal ions for a neutral compound o Roughly equal amounts/concentrations of both acids and base Lecture 37 Titrations - - Neutralize (all acidic protons of) an acid with a base. One concentration is known, the other is not o Titrating an acid of unknown concentration with a base of known concentration o Titrating a base of unknown concentration with an acid of known concentration Used to calculate the total acid or base concentration of a solution - pH only tells you the amount of acid that is dissociated into ions, whereas a titration tells you the full concentration of acidic protons, whether or not they are still attached to the anion Nuclear Chemistry - Nuclide: the nucleus of a specific isotope of an element Nuclear reaction: a reaction that changes the nucleus of an atom, usually causing it to change from one element to another o This is how a lot of the most recent elements were discovered o Different from chemical reactions because chemical reactions involve changes to the distribution of electrons and leave the nucleus alone o Rate of nuclear reactions unaffected by temperature, pressure or catalysts o Energy change for a radioactive reaction can be several million times great than that of a chemical reaction Types of Nuclear Reactions - Gamma radiation: high-energy electromagnetic waves with no change in mass or charge Beta radiation: an electron (e-) emitted as radiation Alpha radiation: a helium nucleus (He2+) emitted as radiation Lecture 38 Radioactive Half Life Half life- the amount of time required for half of a radioactive sample to decay. (concentration of radioactive materials is cut in half)