FRED ZHENG’S SUPER SIMPLIFIED CHEMISTRY FINAL STUDY GUIDE Basics Mass is the amount of matter in the sample -measured in grams Volume is how much room the sample takes up in space -Ways to determine volume -Solids can be submersed in a liquid to see how much liquid it displaces -Liquids can be determined by measuring using a graduated cylinder -Gases have volumes that are equal to the volume of its container Density is a measure of ratio of an object’s m ass to its volume -It is an intrinsic property (one that does not change regardless of amount) of a substance 3 -Density = mass(kg) / volume (m ) Pressure is the force that a sample of gas in a closed container exerts on the container walls or when a solid standing in an environment and the force that a gas is exerting on the walls of the environment -Standard: 760 mmHg = 760 torr = 101.3 kPa = 1 atm Energy is the ability to do work or transfer heat -can exist as heat, light, kinetic energy or chemical bond energy -Kinetic Energy is the energy contained in the movement of molecules -Chemical bonds contain energy -Breaking bonds requires energy; forming bonds release energy -Heat is the transfer of kinetic energy from a body at a higher temperature to a lower temperature -Calorimeter measures Calorimeter measures energy -Temperature is the average kinetic energy of molecules in a sample -Heat Capacity refers to the amount of heat an object must absorb for its temperature to be raised 1 degree celsius -Specific Heat of a substance is the heat capacity of 1 gram of the substance o -Specific Heat of Water= 4.184J/g C
q=mc ᐃT
q=heat m= mass c= specific heat ᐃT= change in temperature -Significant Figures -The amount of digits that are important to ensure accuracy of the report -All non-zero numbers are important -All zeros in front of non-zero numbers are not important -All zeros in between non-zero numbers are important -Zeros behind non-zero numbers are important if there is a decimal point -Accuracy and Precision - Accuracy Accuracy is how close you can get get to the right answer answer -Precision is how close your answers are to each other -Error is Error is |accepted value- experimental value| -Percent Error is Error is Error/Accepted value -Reaction types -Synthesis A+B-> AB -2 reactants combine to form product -Decomposition AB->A+B -1 reactant splits up into 2 products -Single Replacement AB+C-> CB+ A -One isolated element replaces another element in the c ompound -Whether or not the reaction can occur depends on reactivity of each element -More reactive means it will replace -Double Replacement AB+CD-> AD+ BC -2 aqueous compounds are combined to form 2 products, where one is insoluble and precipitates -No reaction occurs when there is no precipitate or when there is a common ion -Solubility can be determined by these rules: -Most silver, lead and mercury salts are INSOLUBLE except their nitrates and perchlorates -Most hydroxides are INSOLUBLE except those of alkali metals and barium
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-All nitrates and perchlorates are SOLUBLE -All alkali metals and ammonium compounds are SOLUBLE -Combustion CH4+2O2-> 2C2O+ 2H2O -An Hydrocarbon, Hydrocarbon, a compound consisting of hydrogen and carbon, reacts with oxygen to produce carbon dioxide and water vapor in a huge forward reaction Atoms Element is any substance that cannot be broken down into a simpler substance by a chemical reaction -Atom is the smallest particle of an element that still retains the chemical properties of that element Compound is a pure chemical substance consisting of two or more different chemical elements that can be separated into simpler substances by chemical reactions -Molecule is the smallest particle of a compound that still retains its chemical properties Parts of Atom -Protons are inside the nucleus with +1 charge and a mass of about 1 amu -Neutrons are inside the nucleus with a neutral charge and a mass of about 1 amu -Electrons surround the outside of the nucleus with -1 charge and a mass of 0 amu -In an neutral atom, protons= electrons -If a neutral atoms gains or loses a negatively charged electron, it will become an Ion -AnionAnion- gain electron; negatively charged -CationCation- lose electron; positively charged -If two atoms of the same element differ in the number of neutrons in their nuclei, they are said to be Isotopes -Elements in nature exist in many different isotopes and we are usually required to calculate the percent composition -Atomic Weight is the average mass number of all isotopes as they occur on Earth -To calculate it, take each isotope, multiply by the percent composition, and add them together -I.E Carbon-12 84% Carbon 14 16% 12 amu* .84+ 14 amu* .16= 12.32 amu
Chemical Reactions and Stoichiometry Diatomic Molecules consist of 2 atoms of identical or different elements -There are 7 important elements that exist as diatomic molecules :O 2, I2, H2, Cl2, I2, Br 2, and F2 Formula Weight: Calculated by adding the atomic weights of all the atoms in the molecule Empirical Formula: Formula: Shows the ratio of atoms within a molecule in lowest whole numbers (HO) Molecular Formula: Shows the number of atoms within a molecule (H 2O2) Percent Composition: The percent of a substance contained within a molecule Converting from Mass composition to Empirical Formula -Convert from percents percents to grams (11% H-> 11g and 89% O->89 grams) grams) -Convert from grams to moles (11g-> 11 mols H and 79g-> 5.56 mols O) -Divide everything by the lowest number of moles (11/5.56= 2 H and 5.56/5.56= 1 O) -That is Empirical Formula (H2O) 23 Mole: Avogadro’s Mole: Avogadro’s number number of 6.02 x 10 -The conversion factor between an amu and gram -For example a 12 gram sample of C contains 1 mole of C Chemical Reactions
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-The bonds that hold together the atoms that make up the reactants break and the free atoms then form new bonds with one another to form new products C3H8(g) +5O2(g) -> 4H 2O(l) + 3 CO2 Reactants Products -This equation tells us that for every molecule of propane (C 3H8), 5 molecules of oxygen (5O 2) are used up and 4 molecules of water (H 2O) are produced with 3 molecules of carbon dioxide (CO 2) -Balancing the equation is making sure that the total number of atoms of each element must equal the total number on the right Stoichiometry -Using the molar ratios of the chemical reaction to find the answer Some tips: -When stuck convert to moles -Factor label is convenient because if things don’t cancel out, you know you did something wrong -When given mass of both reactants, calculate how much of each you have, how much of each you need and stick to the one where you need to least/ have the most of -When converting from grams to moles, make sure you use the molar mass of the compound, not just an element Thermodynamics Entropy- Simply stated, entropy is disorder -Universe fundamentally prefers low-energy states -Universe prefers higher disorder, therefore positive change in entropy is prefered -Denoted as ᐃS -Solid
ᐃG= ᐃH-TᐃS If ᐃG<0 then the reaction is spontaneous in the forward direction If ᐃG>0 then the reaction is spontaneous in the reverse direction If ᐃG=0 then the reaction is at equilibrium
Heat of FormationFormation- The amount of heat that is released or absorbed when 1 mole of a compound is formed from its elements -Heat of Formation of elements is 0 -Usually they give us an equation and we need to find the ᐃH of the reaction and we are given the heat of formation of each compound
ᐃHf for the whole reaction= ᐃHf (Products)- ᐃHf (Reactants)
-Remember if you flip an equation, the ᐃH is negated -Remember if you multiply the equation by a factor, the ᐃH is also multiplied by the factor -Hess’ Law: If a reaction is carried out in a series of steps, ᐃH for the reaction with be equal to the sum of the enthalpy changes for the individual steps Electron Configurations and Radioactivity General -Quantum mechanics increases our understanding of the atom -Electromagnetic energy is quantized as in for a given frequency of radiation, all possible energies are m ultiples of a certain unit of energy called a quantum -Energy changes not smoothly, but in steps -Electrons circle the nucleus not in orbits like the planets, but in orbitals which are regions in which an electron may be found -A collection of orbitals roughly constitutes an energy shell
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-Electrons in higher energy shells are farther away from the nucleus with more energy -There are 4 significant types of orbitals, each with a different shape -Orbitals that have the same shape in a given energy shell comprise a subshell -An s subshell always consists of one spherical orbital -An p subshell always consists of three dumbbell shaped orbital -An d subshell always consists of five oddly shaped orbital -An f subshell always consists of seven oddly shaped orbital -Heisenberg principle states that it is impossible to know both the position and the momentum of an electron at the same time De Broglie's Hypothesis states that matter can be thought of as having properties of both a particle and a wave Electron Configuration
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-Quantum Numbers: Identifies the electron -Principal Quantum Number tells Number tells us which energy shell it is in -Orbital Quantum Number tells Number tells us which subshell it is in -s=0, p=1,d=2, f=3 -Magnetic Quantum Number tells Number tells us which orbital it is in s: 0 p: -1, 0, 1 d: -2,-1,0,1,2 f: -3.-2,-1,0,1,2,3 -3.-2,-1,0,1,2,3 -Spin Quantum Number tells Number tells us what is the direction of electron spin Up: +1/2 Down: -1/2 Aufbau Principle states that a subshell is completely filled before electrons are placed in the next higher subshell Pauli Exclusion Principle states that no 2 electrons in the same orbital can have same spin Hund’s Rule states that electrons would rather be in different orbitals then to be in the same orbital and would only go into the same orbital if no other orbital in that subshell is empty -To write electron configuration of an electron, you have to write out all the electrons before it For each electron: - The number tells us the energy level - The letter tells us the sublevel -The superscript tells us which electron it is in the subshell
Quantum Number of last electron in Oxygen: 2,1,-1, -1/2 2 2 4 Electron configuration is 1s 2s 2p -Valence Electrons are electrons that are in the atom’s outermost shell -Elements with stable octets (8 electrons) are Noble Gases/ Inert Gases Radioactivity and Half Lives -Alpha Decay α -occurs when the nucleus is too heavy -an alpha particle is emitted -made up of 2 protons and 2 neutrons -atomic number decreases by 2; the mass number decreases by 4 -Beta Decay β
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-occurs when there is too many neutrons and not enough protons -a beta particle is emitted -made up of an electron -a neutron is converted into proton -atomic number increases by 1; the mass number stays the same -Positron Emission -occurs when there is too many protons and not enough neutrons -a positron particle is emitted -positively charged electron -a proton is converted into a neutron -atomic number decreases by 1; the mass number stays the same -Gamma Decay -occurs when there is too much energy in the nucleus -releases electromagnetic radiation in the form of gamma rays -nucleus is more stable, but otherwise unchanged Half life is the time in which half of the substance decays away -Take the time and divide by half life to find how many half lives passes -Then divide the mass in half in multiples of the amount of half lives Periodic Table -Vertical columns are called groups -Elements in the same column have same number of valence electrons and usually have similar properties -Horizontal rows are called periods -Elements in the same row have electrons in the same energy shell
Akali Metals -Extremely reactive -1 Valence Electron -Shiny, Grayish white metal Alkaline Earth Metals -Pretty reactive -2 Valence Electrons -commonly found on earth Halogens -Extremely reactive -7 Valence Electron -Combine with metals to form Halides Noble Gases -Not reactive -8 Valence Electron Metals -Malleable means they can be easily hammered into sheets -Ductile means they can be formed into wires -good conductors of electricity and heat -tend to lose electrons in a bond Nonmetals:
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-poor conductors of heat and electricity -tend to gain electrons in a bond Semimetals: -characteristics -characteristics of both Periodic Trends Ionization Energy -Energy required to remove an electron from an atom -The more electrons there are in an energy level, the stronger the pull between nucleus and electron -However, as you go higher in energy levels, the electrons are farther away from nucleus and experience shielding in which electrons from previous energy levels block some of the attractive force -Increases across the table, decreases down the table Electronegativity - How much an atom pulls electrons in a bond -Increase across the table, decrease down the table Atomic Radius -The distance from the center of the atom to the edge -The size decrease across across the table because more electrons= more attractive force -The size increases down the table due to addition of energy level Metallic Character -How metallic that element is -Decreases as you move across the table -Increase as you move down the table Chemical Bonding Ionic Bonds -when an atom in a bond gives up 1 or more electrons to the atom it bonds to -generally formed when the elements that form it have high difference in electronegativities -generally formed between metal and nonmetals -the resulting atoms turn into ions -Lattice energy is energy required to completely separate a mole of solid ionic compound into its ions -the attraction between a positive charge and negative charge is electrostatic force Covalent Bond -when two atoms share electrons -generally formed between 2 nonmetals -Polar Covalent Bond is when one of the atoms hogs all of the electrons due to its high electronegativity -results in one side of the atom having slightly positive charge and vice versa Metallic Bond -when two metals bond -the electrons are donated and move freely throughout the sample and referred to as “sea of electrons” -allows for the good conductivity of metals Single, Double, Triple Bonds -Single: 1 pair of electrons -Double: 2 pairs of electrons -Triple: 3 pairs of electrons -More bonds= stronger and smaller difference between the 2 atoms Bond Energies -Used to predict ᐃH -Find out how many of each bond you have and multiply by the bond energy -Remember bonds have different energies depending on if they are single, double or triple -Take the bond energy of reactants and subtract the bond energy of products Molecular Shapes -VESPR- Valence shell electron pair repulsion states that electron pairs like to find the shape in which they are the furthest from each other -Assume first atom is central atom unless it is hydrogen -Use dots to indicate valence electrons -Add valence electrons until all atoms have octets or are satsified
Number of electron pair sites around the central atom
Number of lone pairs around central atom (not shared)
Shape
Examples
4
0
Tetrahedral
CCl4
7
4
1
Trigonal Pyramidal
NH3
4
2
Bent
H2O
3
0
Trigonal Planar
NO3
3
1
Bent
SO2
2
0
Linear
CO2
Molecular Polarity -Diatomics with polar bonds are polar -Diatomics with nonpolar bonds are nonpolar -Other molecules are generally polar -Unless the central atom has no lone pairs and is surrounded by atoms of only one element, then the polar bonds cancel out and the molecule is nonpolar -Example is Methane CH 4 Phases: Gases, Liquids, and Solids -PhasePhase- refers to where a substance is a solid, liquid or gas Gases -Theoretically, all matter becomes gaseous if its temperature exceeds its boiling pont Ideal Gases -Molecules of an ideal gas do not attract or repel each other -Molecules of an ideal gas occupy zero volume -no gas behaves like an ideal gas -Avogadro’s LawLaw- For a given mass of an ideal gas, the volume and amount (moles) of the gas are directly proportional if the temperature and pressure are constant. -Kinetic Molecular TheoryTheory- states that the kinetic energy of a gas molecule increases proportionally with temperature in degrees Kelvin
Ideal Gas Equation PV=nRT P= Pressure in atm V= Volume in liters n= number of moles R= ideal gas constant of 0.0821 -Partial Pressure -If you have a container filled with more than one gas, each gas exerts a pressure -That pressure is proportional to how many m oles of gas you have over the total moles in container Partial Pressures Moles of gas A Partial Pressure of gas A Total moles in container = Total Pressure of container Intermolecular Forces are responsible for many of the physical properties of matter -London Dispersion Force is the temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles -Occurs in all molecules -Is the weakest force -Larger the molecule (more electrons)= more london dispersion force -Dipole-Dipole are attractive forces between the positive end of one polar molecule and the negat ive end of another polar molecule. -Polar molecules have a partial negative end and a partial positive end. -The partially positive end of a polar molecule is attracted to the partially negative end of another. -Hydrogen Bonds- strong attractive force that occur when hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen to acquire a significant amount of positive charge. -Can be thought of as a super powerful dipole-dipole -Only with elements, O, F, or N Solids, Liquids, and Gases
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-The relationship between a substance’s average kinetic energy and the strength of its intermolecular forces is responsible for determining if the substance will be a liquid, solid or gas Solid - A A substance’s intermolecular intermolecular forces forces are much stronger than the average kinetic energy of its molecules -Molecules are restricted in their ability to move about -Strong forces only permit molecules to merely vibrate in place -Gives a solid a definite size and shape Liquid - A A substance’s intermolecular intermolecular forces forces are slightly stronger than the average kinetic energy energy of its molecules -Molecules in a liquid have enough energy to move past each other -Allows for liquid to flow -Still confined within sample Gas -Molecules in gas are so energetic that they easily overcome intermolecular attraction -Gas molecules spread about to fill the volume of whatever container they are in Properties of Crystalline Solids
Crystal Type
Forces
Examples
Melting point
Hardness
Conductivity
Ionic
Electrostatic attraction
LiF, NaCl
High
Hard, Brittle
Only molten or aqueous solution
Covalent Network
Shared Electrons
Diamond
Very High
Very Hard
None
Molecular
H-bonding, dipole dipole, dispersion
H2O, HCl
Low
Soft
None
Metallic
Electrostatic attraction between cation and sea of electrons
Na, Fe
Variable
Malleable
High
Hydrates (Hydrated salt) is an ionic substance in which water molecules bond to the ions in a fixed ratio. -Copper Sulfate Pentahydrate CuSO 4●H2O -To find percent composition, take the weight of hydrate, evaporate water, take the weight again of just the salt and use basic proportions Phase Changes -The condition of being a solid, liquid or gas is being in a particular state particular state or phase or phase -It is dependant on temperature and pressure
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-The first diagram shows the name of phase changes -The second diagram shows the change in temperature as heat is added to a substance -Moving from left to right, heat is added and the temperature rises -After a while, it reaches a first plateau, meaning although even more heat is applied, the temperature does not change -This is due to the fact that the substance is changing phases -It requires energy for substance to break the intermolecular force and change state -The amount of energy that is required to move from solid to liquid phase is heat of fusion -The amount of energy that is required to move from liquid to gas phase is heat of vaporization -The calculations of these problems go hand in hand with enthalpy -To calculate how much temperature change has occurred with a given heat, find out the temperature that phase change occurs at -Calculate the amount of temperature change that would occur using specific heat -Make sure that if you hit a phase change, you use the heat of fusion/vaporization fusion/vaporization to calculate how much energy is used to convert from one phase to another -Continue on with adding temperature change using specific heat Phase Change in with pressure -Under higher pressure, it is harder for solids to melt and li quids to vaporize -Reduced pressure makes it easier for solids to m elt and liquids to vaporize -This is due to the fact that the pressure is pressing the m olecules closer together and it prevents the molecules from moving around as much and vice versa ***One major exception is water*** -The diagram on the next page shows relationship among temperature and pressure -It shows what state the substance will be when subjected to that pressure and temperature -Anything on the solid line means it is in between 2 states -The Triple Point is the point at which the substance can exist as all 3 states -Critical Point is the temperature and pressure at which the distinction between between liquid and gas can no l onger be made.
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Vapor Pressure -Since temperature is an average kinetic energy, there will always be a couple molecules where the kinetic energy is well above the phase change point -For liquids, a couple molecules have enough energy to escape into the atmosphere through evaporation -Vapor pressure is the pressure that gaseous substance exerts in a closed environment when liquid substance evaporates to gas Factors Effecting vapor pressure -Intermolecular forces: Stronger the forces, the lower the vapor pressure -Temperature: Higher temperature, the higher the vapor pressure -Pressure: Higher the pressure, the lower the vapor pressure Solutions -A solution is a mixture of a solute (the substance which there is less) and a solvent (that of which there is more) -The most commonly used unit for concentration is Molarity -The symbol is M -molarity(M)= number of moles of solute numbers of liters of solution -Molality is also a fairly common unit for concentration -The symbol is m -molality(m) = number of moles of solute number of kilograms solvent -Solubility is how easy a solute dissolves in a solvent -Saturation is the point at which no more solute can be dissolved into the solvent Things that affect solubility -Temperature: more soluble in higher temperatures -Pressure: Only for gases, more soluble in higher pressures -Polarity: Like substances substances dissolve in each other other -Polar molecules will dissolve in polar solvents -Nonpolar molecules will dissolve in nonpolar solvents -When an ionic substance dissolves in water, its bonds break and ions are released into solution and this is called dissociation -Dissociation always creates an equal number of positive and negative charges in solution -The presence of charged particles ables the solution to conduct electricity -Also called Electrolytic Electrolytic solutions and the ions are called electrolytes
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Boiling Point Elevation and Freezing Point Depression -When a solute is dissolved in a liquid solvent, the solvent's boiling point is raised and its freezing point is lowered by a factor ᐃT= kmi k= constant that depends on solvent m= molality i= whole number equaling number of particles in solution
Kinetics and Equilibrium -Kinetics is the study of the rate at which reactant molecules are converted to products in a chemical reaction
-Reactants have to break existing bonds and form new ones to become products -To do so, the molecules have to collide with sufficient energy and proper orientation -Activated Complex is an extremely unstable, high energy arrangement of atoms where the reaction can go either direction since the new and bonds and old bonds are both present Factors that affect reaction rate: -Concentration: With a higher concentration, more molecules are in the same space, resulting in more molecules colliding and reacting faster -Surface Area: More surface area means more contact between the reactants which results in higher reaction rate -Temperature: Increases Increases rate of reactant collisions and the energy involved. More reactants will collide with enough energy to react -Catalyst: Catalyst increase the rate of a chemical reaction without being consumed by it by creating another pathway that has a lower activation lower activation energy. -activation energy is minimum energy needed for the reaction to occur -catalysts are not consumed in a reaction -since catalysts increase the reaction rate in both directions, it does not change the equilibrium -Calculating reaction rate is based on experimental data, you are given a chart with different concentrations and rates x y Initial Rate= K[A] [B] -First find out the order of each reactant by looking at if you change the concentration of it, how much will it affect the rate. -If it doubles when you double concentration, it is first order -If it quadruples when you double concentration, it is second order -If it does not change when you double concentration, it is zero order -Takes each order and plug in values from one experiment from the table and solve for the rate constant -Equilibrium is defined as the point in a chemical reaction at which the concentration of all the reactants and products cease to change -Note* THEY DO NOT HAVE TO BE EQUAL TO EACH OTHER -Many reactions are reversible -Equilibrium expression is only for gas and aqueous solutions aA +bB <-> cC+ dD c d -The equilibrium expresion is Keq= [c] [D] a b [A] [B] -when the forward reaction rate is the same as the reverse reaction rate, we are at equilibrium -Keq= products when K eq> 1 forward reaction is favored
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reactants when K eq<1 reverse reaction favored -For calculations with equilibrium use the SRE or ICE box -Write out equation -Draw Start, React, Equilibrium box -Use the K eq to find out concentration of compounds at equilibrium in conjunction with the SRE box http://www.youtube.com/w http://www .youtube.com/watch?v=IYzBlvUde8 atch?v=IYzBlvUde8U U [Further Instructions] Le Chatelier’s Principle -If some stress is place on a reaction at equilibrium, then the equilibrium will shift in a direction that relieves the stress A+ B<->C -If concentration of A was increased, B would go down and C would go up to make the equilibrium normal again -If concentration of C was increased, A and B would go up to compensate Changing concentration will NOT change K eq
-Heat: The system will shift towards whatever relieves the stress H+I+ Heat <-> J+K If heat was added to the system, the system will shift to the endothermic side If heat was removed, the system will shift to the exothermic side Changing the temperature WILL change K eq -Pressure: The number of particles are important 2A+2B->3C If pressure is added to the system, it will shift toward the side with fewer molecules If pressure is removed from the system, it will shift towards the side with more molecules Equilibrium in Precipitation Reactions -Some combination of cation and anion form very low solubility precipitates. precipitates. The K sp or the Solubility Product Constant is the equilibrium expression for the dissociation of this precipitate -Used it like you would for any equilibrium problem Acid and Base -Water is not all H 2O -Autoionization occurs in which water spontaneously dissociates into OH- an H+ -The equilibrium expression for water is K w= [H+][OH-] -7 -For pure water at 25 degrees celsius, both [H+] and [OH-] = 10 M -p is the abbreviation for -log -pH is -log of concentration of [H+] and it will be 7 in pure water Arrhenius - Acids Acids produce H+ in solution -Bases produce OH- in solution Lewis - Acids Acids are electron electron acceptors in solution solution -Bases are electron donors in solution Bronsted-Lowry - Acids Acids are proton donors in solution -Bases are proton acceptors in solution -Some molecules can act as both acids and bases and are called amphoteric molecules -Acids have pH of below 7 and turn litmus paper red -Bases have pH of above 7 and turn litmus paper blue -Strong acids and bases dissociate COMPLETELY Strong Acids -HCl hydrochloric acid -HBr hydrobromic acid -HI hydroiodic acid -HNO 3 nitric acid -H2SO4 sulfuric acid -HClO4 perchloric acid Strong Bases -Group 1 hydroxides -Weak acids and bases have equilibrium constants for their dissociation K a for acids and K b for bases -Use normal Ice box and equilibrium skills to work with these problems
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-Conjugate pairs are molecules that have identical molecular formulas except that one of them has an addition H+ -Acids produce conjugate bases and bases produce conjugate acids Buffers -Solutions used to minimize a change in pH when an additional acid or base is introduced into the solution -Buffers are made out of conjugate weak acids and bases Titration -Experimental technique used to find out the concentration of an acid by neutralizing the unknown acid or base with a known solution called Titrant -The point at which just enough titrant is added to completely neutralize the the solution is called equivalence point Salt Hydrolysis -When combining equal amounts of strong acids and bases, you neutralize the system as the acid and bases form to make water and salt -However, weak acids combining with strong bases (or vice versa) create salts that will hydrolyze water -The conjugate pair of the weak acid/base will be strong and would react with water to form a basic or acidic solution -Strong Acid + Weak Base= Acidic -Strong Base + Weak Acid= Basic -Strong Acid+ Strong Base= Neutral -Weak Acid+ Weak Base= Base= Neutral Redox and Electrochemistry -Redox are reactions that occur where electrons are transferred from one substance to another -Losing electrons= Oxidation -Gaining electrons = Reduction Oxidation Numbers are a bookkeeping method to keep track of what is being oxidized or reduced 1. -The oxidation oxidation number of an atom in the elemental state is zero. zero. Example: Cl2 and Al both are 0 2. The oxidation number of a monatomic ion is equal to its charge. Example: In the compound NaCl, the sodium has an oxidation number of 1+ and the chlorine is 1 -. 3. The algebraic sum of the oxidation oxidation numbers in the formula of a compound compound is zero. zero. Example: the oxidation numbers in the NaCl above add up to 0 4. The oxidation number of hydrogen hydrogen in a compound is 1+, except when hydrogen hydrogen forms compounds compounds called hydrides with active metals, and then it is 1-. Examples: H is 1+ in H 2O, but 1- in NaH (sodium hydride). 5. The oxidation number of oxygen oxygen in a compound is 2-, except in peroxides when it is 1-, and when combined combined with fluorine. fluorine. Then it is 2+. Example: In H2O the oxygen is 2-, in H 2O2 it is 1-. 6. The algebraic sum of the oxidation numbers in the formula for a polyatomic ion is equal to the charge on that ion. 2Example: in the sulfate ion, SO 4 , the oxidation numbers of the sulfur and the oxygens add up to 2-. The oxygens are 2- each, and the sulfur is 6+. Redox Reactions, Reactions, as reactants from products, one or more atoms are reduced while one or more other atoms are oxidized. -If one takes place, the other must also take place -Just one part of the reaction is the half-reaction -Oxidizing agent causes another species to be oxidized by undergoing reduction -Reducing agent causes another species to be reduced by undergoing oxidation How to balance: balance: 1. Split the equation into two half reactions. 2. Balance any atom other than O or H first. 3. Balance O's using H 2O. Add as many waters as needed to balance the oxygens on the side deficient in oxygen. + + 4. Balance H's using H . Add as many H ions as needed to the side deficient in H. 5. The mass should now be balanced. To balance the charge, determine the charge on each side of the reaction, and add as many electrons to the more positive side as needed to make the charge the same on both sides. 6. Repeat steps 1 through 4 for both half reactions. 7. If the number of electrons lost does not equal the number of electrons gained, find the common multiple and m ultiply each half reaction by the necessary factor. + 8. Sum both half reactions to obtain the balanced, net ionic reaction. In many cases, look for H ions and H 2O molecules to cancel. 9. Check the final equation for mass and charge balance. Electrochemistry - An An electrochemical electrochemical cell (Voltaic (Voltaic or Galvanic Cell) is a device used to produce electrical current from a spontaneous redox reaction -To find if something is spontaneous in electrochem, you need to find standard electrode potential which is the electrical potential difference between the 2 half reactions
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-Look up on the chart the reduction potential and oxidation potential (remember to flip this one if you are using a reduction chart) -Add them to get the overall cell potential 0 -If E >0 it is spontaneous O -If E <0 it is nonspontaenous
-A strip of metal called an electrode is place in each solution -The electrode in the vessel in which oxidation occurs is called the anode and it is negative electrode -The electrode in the vessel in which reduction occurs is called the cathode and it is positive electrode -The Salt Bridge contains a cation and anion that don’t participate in reaction -The salt bridge ions from from to each side to neutralize the reaction and to allow it to continue -Without the salt bridge the buildup of charge from electrons moving from anode to cathode would cause the cathode being too negative to accept any m ore electrons Electrolysis is the process by with electrical energy is put into a nonspontaneous redox reaction to force it to occur -Anode becomes positive electrode -Cathode becomes negative electrode Organic Chemistry -Carbon compounds are important because all living things on Earth are made up carbon -Organic compounds are much more soluble in nonpolar solvents than polar solvents -Organic compounds don’t dissociate in solution HydrocarbonsHydrocarbons- contain only hydrogen and carbon -Alkanes: Alkanes: All single carbon-carbon bonds and end in -ane -Alkenes Contains carbon-carbon double bonds and end in -ene -Alkynes Contains carbon-carbon double bonds and end in -yne -Substituents are an atom or group of atoms substituted in place of a hydrogen atom on the parent chain of a hydrocarbon -Prefixes: -Meth1 -Eth2 -Prop3 -But4 -Pent5 -Hex6 How to name: 1. Find and name name the longest continuous continuous carbon carbon chain. 2. Identify and name groups attached to this chain. 3. Number the chain chain consecutively, starting at the end nearest a substituent group. 4. Designate the location of each substituent substituent group by an an appropriate number number and name. 5. Assemble the name, listing listing groups in alphabetical order using using the full name (e.g. cyclopropyl cyclopropyl before isobutyl). The prefixes di, tri, tetra etc., used to designate several groups of the same kind, are not considered when alphabetizing
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