Second Semester Study Guide Significant Figures (digits) highlights are significant digits - When calculating, try to be as precise as possible - All #’s that are sig. digs. - ALL NON ZERO DIGITS (1,2,3,4,5,6,7,8,9) - ANY ZEROS BETWEEN NON ZEROS (1230004) (123 0004) - ALL ZEROS TO RIGHT OF DECIMAL PLACE - Only if in between non zero numbers or at end of non zero #’s - 1.234567890 - ZEROS BEFORE DECIMAL PLACE ONLY IF THERE IS A DECIMAL - 600, the zeros are non significant (they “fall off the cliff”) - 600. are significant because there is a decimal - Non Significance - ZEROS IMMEDIATELY AFTER DECIMAL PLACE - 234050.00678 (zeros before 678 are NOT SIGNIFICANT) - Scientific Notation Significant Figures - 1.0456 x 1023 = Look only at the first part (1.0456), other part doesn’t matter - Significant Digits while Multiplying/Dividing Multiplying/Dividing - USE VALUE WITH THE LEAST SIG. DIGS. - 6.05204 x 12340.005 = since the first value has only 5 sig. digs, we use that amount of sig. digs. For the answer - Significant Digits while Adding/Subtracting Adding/Subtracting - USE LEAST NUMBER OF DECIMAL PLACES - may include numbers with no decimal places - 19.30451 + 18.4567863 = since the first value only has 5 decimal places, we use that amount in calculation the answer - 1 + 2.5 = since the first value has NO decimal place, we use do decimal places in the value - 1 + 2.5 = 3.5, 3 .5, we round UP to 4 - Rounding Rules - If decimal is between 1 and 4, round down - If decimal is between 5 and 9, round up Density - Density is derive from mass and volume = D= M/V - It’s an intensive property that we derive from two EXTENSIVE PROPERTIES PROPERTIES Atoms and Avogadro’s’ Number - Avogadro’s’ number of particles in an atom = 6.02 x 1023 - Atomic number = number of protons or electrons in an atom (same for protons and electrons) electrons) - Atomic Mass of an atom = neutrons + protons Number of neutrons in an atom = Atomic Mass – Protons Shorthand Diagram (not that pictured on the left) - Mass Number (N + P) over Atomic Number (P, E) - With atomic symbol on the right Isotopes = varying number of neutrons Relative Abundance to Average Atomic Mass - To calculate relative atomic mass = Relative Abundance x Atomic Mass - To calculate relative ABUNDANCE, ABUNDANCE, move % abundance (she will give us this info.) two places to the left (98.97 abundance abundance = 0.9897relative 0.9897relative abundance) - For relative abundance, round to 5 SIGNIFICANT DIGITS
- Once you found the relative abundance, abundance, multiply it to the atomic mass of the isotope that you are dealing with (she will give us this info.) to find relative atomic mass - To find average atomic mass, add up ALL the relative masses and that is answer
-
DON’T FORGET SIGNIFICANT DIGITS AND AMU
AMU = Atomic Mass Unit - 1/12 the mass of a carbon atom is = to 1 AMU - When calculating average atomic mass, DON’T FORGET TO UNIT AMU Nomenclature - Compounds - Ionic Compounds - Polyatomic or Binary Ionic - Molecular Compounds - Cations = metals (positive) - Anions = nonmetals (negative) - Metal + Nonmetal = Ionic Compound - Nonmetal + Nonmetal = Molecular Compound
Chemistry Rules on Naming Compounds Naming Ionic Compounds: - The Metal always goes first in the compound - We only use –ate, or –ite when we are discussing Polyatomic compounds Chloride = in ionic/molecular, the nonmetal ends in – - MgCl2 = Magnesium Chloride ide Naming Molecular Compounds: - Use prefixes - Only use the subscripts to determine, don’t swap subscripts - F2Cl3 = Diflourine Trichloride - S2O = Disulfur Monoxide = in molecular, the first element named keeps its regular name Naming Transition Metal Compounds: - Start with looking at the NON METAL - Determine what subscripts subscripts will have to be used to make them equal - Fe2O3 = Start with oxygen (nonmetal) and determine charge (-6) Then determine determine how Fe2 can match that (add (add a 3+) o o Use the number needed to match and put it in parentheses starting with the transition metal Iron (III) Oxide Then end the equation with the non metal o - If working backwards, start with Transition metal and look to parentheses for its charge Fe3 o Then determine the charge of the nonmetal O2 o o Then swap subscripts Moles
Molecular Mass = Atomic Number of any element - If more than one element, add up the atomic numbers - If subscripts are included, multiply the atomic number by the number subscript - O5 = Molecular Mass = 16 x 5 = 80 Moles to Grams = Number of Moles (will be given) x Molecular Mass (of substance you are trying to achieve) Grams to Moles = Grams (will be given) / Molecular Mass - BE SURE TO ALWAYS USE UNITS AND LABEL WITH GRAMS, LITERS etc. Moles to Volume of Gas @ STP = Number of Moles x 22.4 L (gas at STP) Volume to Moles = Volume of Gas (will be given) / 22.4 L Moles to Representative Particles Particles = Number of Moles x 6.02 x 10*23 Representative Representative Particles to Moles = Number of representative representative particles / 6.02 x 10*23 ** Make a kitchen on the right hand side of paper for calculations, calculations, and LABEL UNITS
Percent Composition = Mass of an element in one mole of that compound / Mass of whole compound x 100 1) Given the total Compound, find its molecular mass 2) Then find each individual molecular mass and divide into big mass 3) Multiply answer by 100 and then you have percent comp * Make sure to remember sig. digs* - When multiplying or dividing, use the least # of sig digs in either value for the final answer
Empirical Formula - Given the Percent Composition, find the lowest possible whole ratio of atoms in a
compound (a.k.a empirical form)
1) Given your Percent Composition, find each of the atom’s molecular mass i.e. = Carbon = 59% / 12.01 which is it’s molecular mass 2) When you have done that to each of the elements and their percents listed, divide the numbers by the SMALLEST ONE - MAKE SURE TO REMAIN CONSTANT WITH NUMBERS OF DIGITS 3) When you have divided by the smallest, use common sense to round up or down to the NEAREST WHOLE RATIO 4) Once you have got whole number ratios, plug that into a chemical formula, as shown above, and that is your Empirical Formula Formula Molecular Formula
1) Given the new percents as shown in the explanation to Number 8, divide by the singular molecular mass, just as Empirical Formula 2) When you have divided by the Molecular Mass, divide each new number by the SMALLEST 3) Using common sense, round up or down to the NEAREST WHOLE RATIO 4) After you have your whole number ratios, find the TOTAL MOLECULAR MASS and divide its molecular mass by it - Simply, take the first set of WHOLE NUMBER RATIOS (5, 7, 1) and set them in a chemical formula C5H7N - Then multiply the ratios (5, 7, 1) by the singular molecular mass, mass, as shown above - 5 x 12.01 - 7 x 1.01 - 1x 14.01 - Then add up ^ to get 81.13 - Then, using the 162.0 g/mol in the given, set up an equation = 162.0/81.13 to get like 1.9967 - HERE, USE COMMON SENSE AGAIN TO ROUND UP TO WHOLE RATIO = 2!! BECAUSE 1.9967 IS CLOSE TO 2 - Once you have that whole ratio, 2, MULTIPLY it to the original equation, C5H7N, so multiply 2 to each subscript - Then the New equation is the molecular formula (C10H14N2) Chemical Equations - Skeletal Equation = Regular Equation that is not balanced - Balanced Equation = Same number of atoms and moles on each side of the equation - Equation Synthesis = LEFT SIDE = REACTANTS RIGHT SIDE = PRODUCTS - Combination Reaction = 2 or more substances reacting or combining to form a new substance
- Combustion Reaction = something reacts with air to produce CO2 and H2O - Include use of a hydrocarbon: hydrocarbon: C?H? , with the ?’s being many different subscripts - Decomposition = single compounds break down into 2 or more simpler products - Single Replacement Reaction = an element replaces another element in a compound - Double Replacement Reaction = 2 ionic compounds change partners in a reaction - Usually occurs in an aqueous solution - Often water is formed Stoichiometry - Mole Ratio = Identifying substances substances in the form of a ratio - Solving Stoich problems: 1) BALANCE EQUATION 2) Put the “known” into train tracks 3) Convert grams into moles if it already isn’t 4) Apply mole ratio 5) Make sure answer uses CORRECT AMOUNT OF SIG DIGS
Limiting Reagent Problem - Limiting reagent either involves comparing of two reactants or products 1) When you see the equation of which you are to find the limiting reagent, balance it 2) Once you balance it, Dr. Burant will have listed the two elements that you are to compare, and their amount in grams/liters as well 3) With those two elements and their givens, calculate the molecular mass and divide by it to find moles 32.46 grams of C x 1 mol of C ------------12.01 grams of C
= 2.703 moles of C
14 grams of O x 1 mol of O --------------32.00 grams of O
= 0.4375 moles of O
4) Once you have completed finding the moles for BOTH (I only did one) reactant, compare each using a mole to mole ratio, and say the BALANCED EQUATION = 2C + 2O2 2CO2 - Once you have a BALANCED EQUATION, EQUATION, look at the two elements that are reacting, here say Carbon and Oxygen, and with their moles found, insert mole ratio to find mole of the other
2.703 moles of C x 2 moles O2 ------------2 moles C
= 2.703 moles of O2
Have( first moles)
Need( second
O2
0.4375
2.703
C
2.703
0.4375
moles)
0.4375 moles of O2 x 2 moles of C ---------------2 moles of O2
= 0.875 moles of C
5) Once you have done all that math, to determine the limiting reagent, whichever element NEEDS MORE THAN THAT ELEMENT HAS is the limiting reagent - So, in this case, O2 needs 2.703 moles, but only has 0.4375 so it is the L.R Percent Yield - This is simple, both the actual yield and the theoretical yield will be given to you Actual Yield x 100 = Percent Yield of that that element element or compound compound -------------- Theoretical Yield Theoretical Yield 1) To find theoretical yield, you must look back at the limiting reagent problem to find the masses given 32.46 grams of C and 14 grams of O2 2) From here, divide by molecular mass, as for L.R, the apply mole ratio, then MULTIPLY by molecular mass of the product that you are comparing it to
32.4 32.46 6 gram grams s of C x 1 mol mol of C mass) = 86.488 grams O2 ------------12.01 grams of C
x 2 mole moles s of O2 x
32.0 32.00 0 gram grams s o2 (mol (molec ecul ular ar
----------------- --------------------------------------------2 moles of C 1 mole of 02
14 grams of O2 x 1 mol of O2 x 2 moles of C x 12.01 grams C (molecular mass) = 5.2544 grams of C -------------------------------------------------------------------------------------------------------------------------32.00 grams grams of O2 2 moles moles of o2 1 mole of of C
3) To decide whichever is the Theoretical Yield, CHOOSE THE SMALLEST OF THE TWO The smallest is 5.2544 grams of C = 5.3 grams of C BECAUSE OF SIGNIFICANT DIGITS 4) When doing either % yield, T yield of L.R PAY ATTENTION TO SIGNIFICANT DIGITS Particles of Gases can be described as: Small particles, Hard Spheres, no significant volume, no forces between them Motion of Gas particles: - Rapid - Constant - Random
3 Variables of Gasses and their Gas Law Problems - Pressure = P - Volume = V - Temperature = T - To indicate which are the first and second values, add the subscript appropriate for that value - Basically, to equate the “before and after” of these values, set both equal to each other in a ratio 1) Make a “garage” in which to specify the known and unknown 2) Convert each to the correct units, to make sure all the units on both sides are the same 3) Write the correct equation in the form of a ratio, with proper subscripts 4) Find the unknown - 3 Types of Variable Equations: - P1V1 = P2V2 - V1/T1 = V2/T2 - P1/T1 = P2/T2
This above is an example of a V1/T1 = V2/T2 - As volume goes up, so does
Direct Graph: temperature This above is an example of a Direct Graph container, when water heats up P1/T1 = P2/T2 - As pressure goes up, so does temperature - In a closed container, when water heats up
- In a closed
This to the left in an inverse graph: P1V1 = P2V2 - In a closed container, that of a vacuum - When volume goes up, pressure goes down Units of Conversion Factors - kPa (kilopascal) - atm (atmospheric pressure) - mmHg (millimeters Mercury) - 1 atm = 101.3 kPa = 760 mmHg Temperature Conversion Units - C = Celsius - K = Kelvin - The temperature temperature of Kelvin is directly proportional to average kinetic energy - Celsius to Kelvin, C + 273 2 73 = Kelvin - Kelvin to Celsius, K – 273 2 73 = Celsius Factors That Affect Gas Pressure - (n) = Molecules - (T) = Temperature Temperature in Kelvin - (V) = Volume in Liters Variation Graphs - As temperature rises, so does the pressure as the molecules are hitting each other more often
- As volume goes down, pressure rises - As the concentration rises, so does pressure General Steps to Solving Gas Law Problems 1) Go to the garage and make a list of known and unknowns 2) Convert units to correct (Kelvin and liters) if necessary 3) Choose the formula to use, depending on variables 4) Predict result and draw a graph 5) Solve; rearrange the formula to solve for unknown 6) Make sure answer is correct and in correct number of sig digs Different types of gas law problems - Boyles Law = P1V1= P2V2 - Inverse relationship - Charles Law = V1/T1= V2/T2 - Direct relationship (watch for temp, must be in K) - G – L’s Law = P1/T1=P2/T2 - Direct relationship relationship - Combined Gas Law = P1V1/T1= P2V2/T2 - Ideal Gas Law = PV=nRT - Pressure (Volume) = Moles(R) (Temperature) (Temperature) - When finding R, if you are dealing with ATMs, use 0.0821 - If dealing with kPas, use 8.31 - The R means the constant Atomic Radius - Half the distance between the nuclei and two atoms of the same element, when the elements are joined. Periodic Trend - As you move Across, the atoms Decrease - As you move Down, the atoms Increase
Group Trends - The reason size increases as you move down, is because the atoms of the elemen elementt are are furth further er awa away y from from t he nucl nucleus eus and and there there is is less less an an attra attracti ction on (shielding effect) Ionization Energy - The energy required to remove an electron when an element is in a gaseous state (opposite of atomic radius) - Ionization energy Increases as it goes Across, and Decreases as it goes Down - as you go down, the electrons are farther away from the nucleus, and therefore little energy is needed to remove them, because they are not strongly attracted Electron Dot Configuration - What are they? = They are like the address of an element - What do they include? = Energy Level (period number ((1 – 7)), Orbital shape (s, p, f, d) and the number of electrons they contain (subscript ((1 – 8)) - How to write them - Locate the element - Determine the number of valence electrons is has - Find which block (s, d, f, p) it is in - Determine the energy level (1 – 7) - Note the number of electrons in orbital - Work backwards until 1s2
- Big Exceptions = the 3d block is actually in period 4 (4d in five…)
The Green “1s” should in theory go next to the purple “1s”. - Lets start with Neon, which is directly below the "1s" - Neon is in period 2, section P, and has 6 valence electrons = 2p6 - 2d6, (then go backwards to 2s2, then end at 2s1 - 2p6, 2s2, 1s2 Electro negativity (atom gains a slightly negative charge) - The ability of an atom to attract electrons when elements are in a compound (water molecule for example) * Does not affect noble gases because they have all their electrons - As you move across, the electro negativity gets Stronger (increases)
- As you move down, there is less an attraction and the electro
negativity decreases
Atomic Radius goes from Francium up to the noble gases (first arrow) Electro Negativity goes from Fluorine down to the left (second line) Ionization Energy starts at the right and goes down to the left (last arrow) Atomic Spectra - When atoms absorb energy, electrons "jump up" to higher energy levels. When the electrons move back down to their ground state, they emit energy. - Each element has its own unique atomic emission spectra Bonds Four types: ionic, metallic, covalent network, molecular compounds (covalent bonds) Metallic = occurs between positive ions (sea of electrons) - Malleability and ductility - Medium bond strength, variable melting point, conductivity in H2O, as a solid AND a liquid - Co, Cu, Fe Ionic = positive and negative ions (cations and anions) - Strong bond, high melting point, conducts electricity in H2O, Conductivity as a liquid but NOT a solid - NaCl Covalent Network (atoms) - Very strong bond, does not melt, Does Not conduct at all - Diamonds Molecular Compounds (Vander Waals Forces)
- Three types of bonds: Hydrogen (strongest), (strongest), Dipole – Dipole (second), London Dispersion Forces (weakest) Bonding Symbols - -, .. = single bond - - = Double bond - - - = Triple bond How to draw Lewis Dot Structure - Figure out the atom with the central, most bonds (always carbon if dealing with it) - Determine number of bonds the other atoms can form
Here, the central atom is Carbon (it can have up to 4 bonds) -
Chlorine, on the other hand, can only have 1 bond, but there are still 3 other places to fill - To fill the other places, use dots
-
EXCEPTIONS = NO2 - BORON HAS 3 spots - HYDROGEN HAS one
Geometry - Diatomic,
Triatomic =
- Bent Triangular - Trigonal
-
-
(water) Planar
Pyramid
Tetrahedral
VSEPR Theory (Valence Shell Electron Pair Repulsion) - Valence electrons "want" to be as far away from other pairs as possible - This influences the shape of molecules - The "dots" that are apparent in the structure indicate the pairs of unshared pairs (one pair of dots = one unshared pair) - The number of unshared pairs also determines determines the shape of the bond - 4 shared pairs and 0 unshared = Tetrahedral Tetrahedral - 3 shared pairs and 1 unshared = Pyramid - 3 shared pairs and 0 unshared = Trigonal Planar - 2 shared pairs, and at least 1 unshared = Bent Shape (water molecule) * All diatomic shapes are linear Bond Polarity - How equal the bond sharing is - Measured by finding the difference in electro negativity - Any time there is a bond between 2 different atoms, it is polar
Molecular Polarity - One end of the molecule is slightly different than the other - If the molecule is symmetric, it is non polar * The molecule might have polar bonds, but it could still be non polar Hydrates - A compound that incorporates H2O in its solid structure, in a much defined ratio - CaSO4 . 2H2O (for every formula unit of CaSO4 there are 2 water molecules) - Water in a hydrate is called hydration, when it is removed, it is called anhydrous or anhydrate Solutions - Homogenous Solutions are aqueous solutions, that containing water - Solvent = dissolving medium - Solute = stuff that's dissolved - On a micro level, the polarity of H2O attracts ions, not all ionic bonds dissolve in water, "likes dissolve likes" Molarity - Number of moles of solute per unit of solution - M = Moles / Liters *must be in liters
In the problem above, you are to find grams (moles in this case) so you need to transform the equation to solve for moles - Then you need to change mL into Liters, and multiply liters by the molarity to get moles - Finally, find the molecular mass of Epsom and MULTIPLY the MULTIPLY the number of moles by molecular mass Dilution - M1 (V1) = M2 (V2) - The volume may either be used in mL of L, but be consistent Hydrate Problems - With recorded lab data, one should be able to conduct the following
Thermochemistry - The study of energy in chemical reactions, and their change of state - Heat is energy transferred from one object to another because of temperature difference between them - Heat flows from warm to cold - How is heat measured? - Units: calorie = quantity of heat needed to raise temperature of a gram of water 1 degree Celsius - The calories in food are BIG C’s, called kilocalories - Joule: standard unit of energy - What is specific heat? - The amount of heat required to raise the temperature of 1 gram of a substance 1 degree Celsius - Formula for specific heat: - C (specific heat) = q (heat) (heat) x M (mass) x T (temperature) - In solving specific heat, if the temperature goes DOWN, it is ENDOTHEORMIC (Positive Answer) - If it goes UP, it is EXOTHOERMIC (Negative Answer) Hess Law - Allows us to determine change in heat for a reaction by adding all the heats for an intermediate intermediate reaction
Rates of Reactions - Why do reactions occur? - Collisions theory explains that reactions happen when colliding particles have enough kinetic energy - The minimum amount of energy needed for a reaction to occur is called activation energy
Factors that affect the rate of a reaction - Temperature: if the temp is increased, the rate of reaction increases - Greater Concentration: greater concentration increases reaction rate - Particle Size: if you make particles SMALLER, you increase surface area and reaction rate increases - Catalyst: increase rate without being used up Reversible Reactions - Reactions that occur simultaneously in both directions in a closed system * Noted by double arrows Equilibrium = when the rate of a forward reaction and reverse reverse reaction equal each other Lachateliers Principle - If stress is applied to a system in a dynamic equilibrium, the system reacts to counteract the stress - Examples of stress: addition of removing of products/reactants, products/reactants, lowing or raising temperature, temperature , increase or decrease pressure Equilibrium Constant (KEQ) - The ration of products to reactants in a balanced system - For Keq, use only aqueous or gram solutions, no solids or liquids
S2 (g) + 2 H2 (g)
2 H2S (g)
[H2 (g)] = 2.16 mol/l or 2.16 M [S2 (g)] = 0.3 mol/l or 0.3 M [H2S (g)] = 0.5 mol/l or 0.5 M
Answer: Step 1: Write the Keq expression (reactants over products) [H2S (g)]2 Keq = [H2 (g)]
2
*
[S2 (g)]
Step 2; Plug in values: [0.5 M ]2 Keq = [2.16 M]
2
*
[0.3 M]
Keq = 0.179 (the Keq has no units) - If the Keq is greater than one, than the Products are favored, and if it is less than one, the Reactants are favored - Acids and Bases - Acid: pH of 1 – 7, ionized hydrogen atoms, conducts electricity Different Acids: HCl = hydrochloric acid H2SO4 = sulfuric acid HNO3 = nitric acid H3PO4 = phosphoric acid CH3COOH = acetic acid (last H gets removed) H2CO3 = Carbonic acid Bases - cleaning solutions - Slippery feel to them - Hydroxide ions - NH3 = ammonia Arrhenius - He said acids are compounds containing Hydrogen ions that ionize in water - He said bases are compounds containing OH ions that ionize in water Bronsted – Lowry - Acids are hydrogen ion donors - Bases are hydrogen ion acceptors Conjugate Pairs - Two substances that are related, on either side of the reaction - Must be opposite pairs (acid and base vice versa) pH – way of determining acidity or basic concentration concentration in a substance - To solve, use – [ log ] (for sig digs, start counting after the decimal places) (pH + pOH = 14) Strengths and Concentrations of Acids and Bases - Concentration = how much molarity is contained in that solution - Strength = degree to which acid or base ionizes in water - Strong acids = completely ionize - Weak acids = partially ionize Ka = [ H+ ] [ Resulting Acid (without H+ ] / [ Beginning Acid ]
