Understanding Atomic Structure and Chemical Compounds, Exams of Nursing

Download Understanding Atomic Structure and Chemical Compounds and more Exams Nursing in PDF only on Docsity! 2024 CHEM-105 EXAM STUDY GUIDE UPDATE RATED A+ Chemistry ● the study of matter ● transformation of matter, examples of matter include chairs, air, the ocean, and you! ● transformation= change ● If it takes up space, it’s matter! Physical Transformation ● does NOT alter the chemical composition or “identity” of the substance, just the physical state ● ex. ice→liquid→gas Physical Changes ● Change in state (gas, liquid, solid) ● Change in size ● tearing ● breaking ● grinding ● dissolving (ex. salt in water) Remember that the chemical composition remains the same. Chemical Transformation ● alters the chemical composition and identity of the substance (ex. burning wood) Chemical Changes ● generation of gas ● generation of solids (know at precipitates) ● disappearance of a solid ● change in color ● change in temperature September 2, 2015 Chemical Property ● the ability of a substance to combine with or change into one or more substances (ex. iron forming rust when combined with oxygen) Physical Property ● characteristic that can be observed without without changing the sample’s composition (ex. density, color, odor, hardness, melting point) Scientific Method ● set of steps used in scientific inquiry ● we make observations, come up with questions, and create theories 1. Observation: we observe facts (which may or may not be reproduced) ex. salt is white, can split into tiny cubes, melts at high temperatures 2. Hypothesis: guesses or suggestions as a possible explanation of a set of facts ex. my son’s runny nose is due to a cold, my son’s runny nose it due to allergies; both are hypotheses and they can either be proven or disproven through experimentation 3. Theory: an explanation based upon a large number of facts or hypotheses that have been widely accepted due to extensive testing, however, theories can be disproved by more experiments and revert back to hypotheses ex. all matter is made up of tiny, individual particles called atoms. Significant Figures ● number of digits known exactly in a measured quantity, plus one more (which is estimated) Rules for Determining Significant Figures 1. all non zero numbers are significant ex. 53.7 = 3 sig. figs. 2. “in between zeros” (zeros that are between nonzero digits) are significant ex. 1001 = 4 sig. figs. 3. trailing zeros in a number WITH a decimal point are significant ex. 0.20 = 2 sig. figs. 12.000 = 5 sig. figs. The zero to the left of the decimal point is NOT a significant figure. 4. leading zeros are NOT significant ex. 0.004 = 1 sig. fig. 5. trailing zeros in large numbers without decimal points are NOT significant ex. 12,000 = 2 sig. figs. 12,000. = 5 sig. figs. 6.no significant figures in the powers of 10 in exponential notation ex. 3.60 x 10^4 = 3 sig figs All digits before are significant! (m)→10^-3 Micro (u)→10^-6 Nano (n)→10^-9 Factor Label Method ● it’s easy to convert if meanings are clear ex. 24cm = ? km Cucumber example: 4 cukes = 1 dollar (conversion factor) How much are 8 cukes? 8 cukes x (1 dollar/4 cukes) = 8/4 →$2 the unit that is cukes, cancels out ● you may need to use more than one conversion factor September 10, 2015 English to Metric ● 1.6 Km = 1 mile ● 2.54 cm = 1 inch ● 1 Kg = 2.2 lbs Base Units of the Metric System length: defined in terms of meters ● originally a metal bar, now a wavelength of light mass: defined in terms of grams ● mass is the quantity of matter contained in an object ● reference standard is a Kg (that is a Pt/Ir alloy→does not decompose) ● invariant with location→ mass stays the same despite location ● mass is measured with a balance (a comparative device) weight: a measure of gravitational attractive force between two objects ● varies with distance between two bodies (measurements taken at 11,000 ft and at the bottom of the ocean would be very different) ● measured using a scale Temperature Scales Temperature: measures the hotness or coldness of an object, measures using a thermometer ● must have two references and a scale between them ● boiling point and freezing point of water are used as reference points ● Celsius: 0→100 Fahrenheit: 32→212 Converting between fahrenheit and celsius 1. add 40 2. if celsius, multiply by 5/9 if fahrenheit, multiply by 9/5 3. subtract 40 Kelvin Scale: is an absolute scale K = C + 273 C = K - 273 ● the temperature depends upon movements of molecules slower the movement, the colder the temp ● at 0 kelvin, all molecular motion stops Density density= mass/volume ● mass is measured in grams ● volume is measured in mL ● density is g/mL Intensive property: does not depend on the amount of substance measured (ex. color, melting point, temperature) Specific Gravity Specific gravity is equal to the density of the substance divided by the density of water ● has NO units ● density of water= 1.0 g/mL ● used to measure urine and tells us something about the excretory system, can be measured by a hydrometer or laser STATES OF MATTER September 11, 2015 Energy ● the capacity to do work ● what is “used up” when work is performed ● is the “currency” used to perform work, just as the dollar is the currency used to buy goods Energy can be assigned into two, broad categories ● kinetic energy : energy of motion, ex. person walking, speeding cars KE= ½ mv^2 ● potential energy : stored energy, has the potential to do work a. energy of position : the energy of an object that is either attracted or repelled (think of a spring) by some other object, the greater the height→the greater the potential energy b. chemical energy : stored in bonds that connect atoms (holds atoms together), when theses bonds are broken, energy is created, ex. gasoline, food Heat ● the flow of energy from one object to another due to a temperature difference ● “heat energy in transit” ● hit always flows from hot to cold, ex. a cup of coffee left in a room ● mass is conserved in all chemical transformations! , ex. CO + PbO → CO2 + Pb Atoms consist of three subatomic particles 1. protons ● always have a positive charge 2.neutrons ● have the same mass a protons ● have no charge 2. electrons ● lighter than protons and neutrons ● value is not included in the mass of the atom ● have a negative charge Structure of Atoms ● have a specific arrangement ● Nucleus: small, dense, positive charge, located in the center of the atom, contains protons and neutrons, ● electrons surround the nucleus, diffuse region of negative charge Simple Model ● commonly called the Bohr model ● electrons move around nucleus in orbitals ● orbitals have a fixed energy values called quantum levels The Atomic Symbol ● atomic mass: total protons and neutrons ● atomic number: number of protons or electrons ● charge: + or - values Isotope ● an element with the same number of protons, but a different of neutrons ● atomic number does not change, but the atomic mass changes ● subtract atomic number from atomic mass to find the number of neutrons ● isotopes of various elements consist in nature in varying percentages September 16, 2015 Atomic Mass ● the atomic mass of an element is the “weighted average” of the mass numbers of all the naturally occurring isotopes Periodic Table ● a way of organizing the elements ● Mendeleev organized the elements based on both their atomic mass and their properties ● Law of Mendeleev: properties of the elements recur in regular cycles periodically when the elements are arranged in order of increasing atomic number ● Periodicity: repetitive properties of elements ● the columns of the periodic table are know as groups ● the rows of the periodic table are known as periods ● Group A: representative elements, main group elements ● Group B: transition elements ● There are also inner transition elements Electrons ● are distributed in discrete shells that surround the nucleus ● the shells are designated by the principle energy level or “quantum number” ● maximum capacity can be calculated by 2n^2, where “n” is the shell number ● within shells, we have subshells ● each subshell contains orbitals and each orbitals can ONLY contain two electrons Electron Configuration ● distribution of electrons within an atom ● Three Rules: 1. orbitals are filled in order of increasing energy 2. each orbital can only hold two electrons, and must have opposite spin (opposite directions)→Pauli Exclusion Principle 3. in the p subshell, you only put 1 electron in each orbital until the orbitals each have one, then you can pair them ● can be written in an expanded or condensed format ● the noble gas configuration can help to shorten electron configurations for example, you can replace 1s2 with He Orbital Box Diagrams ● a box represents an orbital ● an arrow represents an electron ● a pair of arrows with heads in the opposite directions represents a pair of electrons with paired spins Lewis Dot Structure ● the symbol of the element represents the nucleus and filled shells ● dots represents valence electrons ● put one electron around each box, then pair them up (different from chem lab) Ionization Energy ● the energy needed to remove a valence electron ● increases as you go left to right on the periodic table, and also as you move up the periodic table ● as you go across, the valence electrons are in the same shell and subject to increasing attraction as the number of protons in the nucleus increases ● as you move up, the valence electrons are lower in principle energy levels, which are closer to the nucleus and feel the nuclear charge more strongly (have fewer shells), the farther away the electrons are from the nucleus, the easier it is to remove them ● when a neutral atom gets a charge, it becomes an ion ● metals tend to lose electrons Chapter 3 Two main types of compounds Molecular ● basic smallest particles the molecule ● two or more atoms bonded together by electron sharing (covalent bond) Ionic ● compounds in which the smallest particles are electrically charged irons ● held together in lattices by electrostatic attraction (between metal and nonmetal) Rules for Naming Binary Molecular Compounds 1. less electronegative element goes first using prefixes (di, tri, tetra, penta) 2. second element Drawing Lewis Structures of Covalent Compounds ● aiming to satisfy the octet rule for every atom 1. count up the total number of valence electrons 2. find the central atom→whatever atom appears the least→if they all appear in the same amount and there is a carbon involved, it is always carbon 3. put a shared pair of electrons between the central atoms, and the surrounding atoms (lines= shared pairs dots=unshared pairs) 4. distribute electrons around the surrounding atoms until the octet is filled 5. place remaining electrons around the central atom, and if the central atom does not have 8 electrons, move a pair from an outside atom (make a double bond, triple bond, etc.) Valence Shell Electron Pair Repulsion Model ● predicts the bond angles and 3 dimensional structure of molecules ● determined by number of pairs around the central atom and number or unshared pairs ● September 30, 2015 Shapes ● knowing the shape can predict whether the molecules is polar or nonpolar Chapter 4 Chemical Reaction ● substances combine to form chemically different compounds ● Reactants→Products ● Reactants are substances that react ● Products are substances that are produced Chemical Equation ● an expression in which symbols and formulas are used to represent a chemical reaction ● example: 2H2 + O2→2H2O ● the number that precede the chemical formula are coefficients that tell you how much you have ● whatever you start off with on one side, you must have the same amount on the other side→balance the chemical equation ● the arrow represents “yields” “produces” Balancing Chemical Equation 1. write correct formulas for reactants and products, connected by an arrow. These do not change 2. adjust the coefficients (to the left of the formulas) so that there are an equal number of each element’s atoms on both sides *keep polyatomic ions together, and balance hydrogen and oxygen atoms last/ October 5, 2015 Formula Mass (weight) ● sum of the atomic weights (in amu) of ALL the atoms in a chemical formula ● simply add up the weights of each of the atoms, which are given on the periodic table ● only need to round to the tenths place Molecular Weight ● same as formula mass ● should only be used for covalent compounds Moles ● Avogadro’s number→6.02x10^23 ● this number of hydrogen atoms equals 1.0 grams ● the “molar mass” is the mass of one mole of any substance expressed in grams ● molar mass is equal to its atomic mass in grams ● just the units change ● coefficients in a balanced chemical equation tell how many molecules/moles are needed and how many molecules/moles are formed ● coefficients can be put into ratios to create conversion factors ● think of it like a recipe→doubling/halving, etc. October 7, 2015 Calculations Using Balanced Equations 1. Make sure you have a balanced equation 2. Write down (w/units) what you are given; write units of what you’re seeking 3. Get conversion factors from mole ratios of the equation and from molar masses Another application of the mole concept is in determining how many particles we have in the sample ● grams→moles→Avogadro’s number ● simply find how many moles are in sample (when given grams), then multiply by Avogadro’s number Limiting Reactant ● determines how much product will form ● is used up totally in a reaction ● other reactants are present in excess, and part of the substance remains unreacted Types of Yields ● actual→the mass of the product formed ● theoretical yield→what we get through our calculations, the predicted yield doesn’t take into account any material lost in the reaction ● percent yield→the ratio of the actual and theoretical yield multiplied by 100, specify the efficiency of a particular chemical reaction percent yield= actual/theoretical x 100 October 9, 2015 Ionic Reactions in H2O ● recall, ionic compounds separate in aqueous solutions ● they break up into their ions Reactions between ions occur IF ● 2 ions form an insoluble precipitate ● a gas is released ● acid/base neutralization ● oxidation/reduction of ions occur Precipitation ● occurs when 2 ions come together to form an insoluble salt GO OVER SOLUBILITY CHART the solubility chart does not cover every exception Oxidation -Reduction (REDOX) Reaction ● Oxidation→loss of electrons ● Reductions→gain of electrons LEO says GER ● in a reaction, something must be oxidized and something else must be reduced October 13, 2015 Alternative Definition (Redox) ● Oxidation→gain of oxygen and/or the loss of hydrogen ● Reduction→loss of oxygen and/or the gain of hydrogen ● Redox reactions always occur as a pair ● the species that is oxidized is called the reducing agent since it causes the reduction of the other species by donating its electrons to it ● the species that is reduced is called the oxidizing agent since it causes the oxidation of the other species by accepting its electrons ● always look at the left side of the equation Heat of Reaction ● the amount of heat given off or absorbed in the course of a chemical reaction ● endothermic→means heat is required, absorbed, can be included in balanced equation as a reactant (calories or joules) ● exothermic→heat is released as a product of the reaction (written on the right) ● can be used to solve for energy ● used in the same way as in stoichiometry problems October 14, 2015 Chapter 5 Gases ● least dense state of matter ● mostly empty space between molecules ● molecular motion is very rapid ● expand to fill their containers (diffusion) very quickly ● describe gases with four variables: Pressure (P), Volume (V), Temperature (T, always in kelvin), and amount (n, the number of moles) Pressure ● there are varying degrees of strength, with hydrogen bonding being the strongest and london dispersion force being the weakest D ip o l e - d ip o l e ● between polar molecules ● polar molecules have dipoles→one end is negatively charged, one end is positively charged ● use dotted lines to represent intermolecular attraction Hydrogen Bond ● an especially strong dipole ● results when H is covalently bonded to very electronegative F, O, or N atom ● is what gives water its high boiling point, surface tension, the reason ice floats, among other properties London Dispersion Forces ● between nonpolar molecules ● nonpolar molecules have a symmetrical cloud ● a temporary dipole occurs when an electrons are briefly uneven ● low boiling points Properties of Liquids Surface Tension ● tendency of a liquid to minimize its volume (result of intermolecular forces) ● unbalanced intermolecular forces=bunched molecules=surface tension ● larger intermolecular forces→larger surface tension Evaporation ● molecules near the surface of the liquid gain enough kinetic energy to escape into the gas phase ● the kinetic energy is greater than the intermolecular forces ● known as a vapor ● is an endothermic process→absorbs energy from the surroundings ● the thermal energy required for a liquid to transform into the gas phase is called the heat of vaporization October 23, 2015 ● vapor pressure→the partial pressure exerted by molecules that have escaped into the gas phase ● dynamic equilibrium →when the amount of molecules escaping into the gas phase equals the amount of molecules returning into the liquid phase Boiling Point ● when vapor pressure is equal to the atmospheric pressure ● boiling occurs when the molecules of the vapor can push back on the atmosphere with equal force ● when you change the pressure, you can change the boiling point (lower pressure, lower boiling point) ● affected by intermolecular forces and the size and shape of the molecules ● the stronger the forces, the the higher the boiling point ● CH4 (nonpolar molecule, london dispersion forces, weight =16 amu) has a lower boiling point than H2O (polar molecule, hydrogen bonds, weight=18 amu) ● Non-polar molecules→larger the size the higher the boiling point (larger surface area= larger london dispersion forces) ● Non-polar molecules of similar size (molecular weight), the molecular shape can help to determine the boiling point→larger surface area=larger london dispersion forces ● ex. spheres are easier to break apart than linear molecules Properties of Solids ● if molecules are organized in an orderly fashion the solid is→crystalline ● if the molecules are disorganized the solid is→amorphous (without form) ● generally, a substance in the solid phase is more dense than in the liquid phase (exception is water) October 26, 2015 Crystallizatio n ● is the process of formation of a solid from a liquid ● is an exothermic process ● the thermal energy given off by crystallization is called the heat of fusion Phase Change ● going from one state to another ● ex. liquid to solid ● transition from liquid to gas→heat of vaporization ● heat of fusion of water→80 cal/g ● heat of vaporization of water→540 cal/g ● sublimation→solid to gas Chapter 6 Solutions ● a homogeneous (same throughout) mixture of 2 or more substances in a single phase ● ex. a saline solution (NaCl and H2O→solid dissolved in a liquid) Solvent ● usually a liquid in which other substances are dissolved ● especially water→”a universal solvent” ● have the most of Solute ● substance dissolved in the solvent (could be a gas, liquid, or solid) ● have the least of October 28, 2015 Characteristics ● components won’t separate out ● transparent to light ● can be separated by physical means (ex. distillation), but not by filtration Qualitative Terms ● dilute→means relatively small amount of solute in solvent ● concentrated→means a lot of solute dissolved in the solvent Quantitative Terms ● saturated solution→maximum amount of solute dissolved at a given temperature→addition of more solute, will NOT dissolve ● this represents the solubility values found in tables→there are specific numbers ● unsaturated solution→less than saturated concentration→more solute would dissolve if added ● supersaturated→an unstable condition in which more solute is dissolved than the normal max at a given temperature (can be created by heating up the solution, and cooling it down slowly)→reverts quickly back to saturated solution with seed crystal or agitation Solubilities and Temperature ● solids and liquids are usually more soluble with heating (ex. adding sugar in iced tea vs. hot tea) ● gases are less soluble with heating→the warm the solution, the less gas is in the solution Solubilities and Pressure ● has little effect on solid or liquid solubility ● increasing partial pressure of a gas always increases solubility (Henry’s Law) ex. the “Bends” October 30, 2015 Predicting Solubilities ● “like dissolves like” ● nonpolar solutes best dissolve in nonpolar solvents (like hydrocarbons) ● polar and ionic solutes dissolve best in polar solvents (especially H2O) ● polar and nonpolar do not mix ● a liquid that is soluble in another liquid→miscible, ex. ethanol in H2O ● a liquid that doesn’t dissolve in another liquid→immiscible, ex. oil in water Representing Concentration 1. % concentration 2. molarity 3. parts per million/parts per billion (ppm/ppb)→extremely small! Percent Concentration ● part/whole x 100 a. w/v%→(g solute)/(mL solution) x 100 b. v/v%→(mL solute)/(mL solution) x 100 c. w/w%→(g solute)/(g solution) x 100 ● remember that solution is solute + solvent!! Molarity ● M= (# of moles solute)/(# of liters solution) ● # of moles→grams solute/molar mass ● exothermic reaction →energy of the products is less than the energy of the reactants, energy is released, heat of reaction (energy) is this difference ● endothermic reaction →much slower since activation energy is high and heat is absorbed Factors Influencing Rate A. Nature of reactants ● ex. H2 vs. Fe ● H2→fast, burns ● Fe→much slower, oxidizes ● ionic solutions tend to react faster than molecules B. Temperature→rate ● greater chance of collisions at high temp. ● heating→every time you heat the solution 10°C, the reaction rate doubles ● more molecules have an energy value that is greater activation energy C. Physical State ● smaller particles react faster ● homogenous gas mixes or solutions are faster ● heterogenous reactions are slower D. Concentrations of Reactions ● usually the rate is directly proportional to concentration ● as concentration increases, so does speed of reactions ● the way react concentration is increased depends on the state of the matter involved ● liquids (higher concentration of solute particles) ● gas (high pressure) ● solid (larger surface area, powder reacts faster than chunks) November 13, 2015 E. Catalyst ● substance which speeds reactions without being consumed in the process ● ex. enzymes are proteins that act as biocatalysts ● catalysts supply a lower activation energy pathway for the reaction (a different mechanism ● finds the shortest way to get from point A to point B Reversible Reactions and Equilibrium ● most reactions are irreversible, ex. digesting food, burning paper ● some are reversible, can be made to go in either direction ● two arrows are used to indicate a reversible reaction: ⇄ ● reactants form products, products form reactants→ all four species are present Equilibrium Constant ● for every chemical equilibrium ● aA + bB ⇄ dD + eE ● capital letters are reactants, lower case letters are coefficients ● there is an equilibrium constant→ K K= [D]^d[E]^e/[A]^a[B]^b ● brackets represent molar concentration, just write the the formula unless the concentration are given ● products are always on top ● if K is large there is more products than reactants→ reaction lies to the right ● if K is small there is more reactants than products (K<1) → reaction lies to the left ● K will be constant at a given temperature ● if something is at equilibrium, there is no net change in the concentration ● Note: K’s tell how far a reaction goes to to products or equilibrium, while rates tell us how fast the reaction gets there THE TWO ARE NOT RELATED L e C h a t e ll i e r ’ s P r i n c ip l e ● if you apply a small stress to the system, the system shift direction to return to equilibrium ● if a system in a state of dynamic equilibrium is subjected to an external stress (change in concentration, T, P), the equilibrium will shift in such a way to remove the stress and return to the state of equilibrium Haber Process ● N₂ (g)+ 3H₂ (g) ⇄ 2NH₃(g) ● if more N2 is added→reaction shifts to the right to use up some of the nitrogen ● if more H2 is added→reaction shifts to the right to use use some of the hydrogen ● is some H2 is removed→reaction shifts to the left to replenish the hydrogen ● some NH3 is removed→reaction shifts to the right to replenish the NH3 November 16, 2015 Effect of Temperature ● an exothermic reaction releases heat and behaves as follows when the temperature is changed ● A + B ⇄ D + E + heat ● what happens to this equilibrium if you raise it’s T (add heat)?→move to the left ● if you cool it?→move to the right ● an endothermic reaction absorbs heat and behaves as follows when the temperature is changed ● A + B + heat ⇄ D + E ● what happens to this equilibrium if you raise it’s T (add heat)?→move to the right ● if you cool it?→move to the left Effect of Pressure ● N₂ (g)+ 3H₂ (g) ⇄ 2NH₃(g) ● pressure changes only affect equilibrium involving gases ● look at the number of gas molecules, the fewer the gas molecules, the less pressure (4 and 2) ● increase the P→move to the right ● if you reduce P→move to the left ● if the number of molecules is the same on both sides and you apply pressure, nothing happens Effect of a Catalyst ● has no effect on the final position of an equilibrium, only how quickly the equilibrium is reached Chapter 8 Arrhenius definitions ● acid→a substance which increases hydronium ions [H₃O+] in aqueous solutions ● base→a substance which increases hydroxide ions [OH-] in aqueous solutions ● only pertains when you put something in water Learn Strong Acids and Bases (table 8.1) Acids ● HNO3, HCl, H2SO4, HBr, HI, HClO4 Bases ● NaOH, KOH, LiOH, Ba(OH2) Weak Acids ● don’t react completely with H2O, mostly HA in solution, only a little (say 1%), H₃O+ added ● usually have equilibrium symbol ● HA + H₂O ⇄ H₃O+ + A- ● H in HA is the acidic hydrogen, donates a proton to water ● a proton donor is an acid, and a proton acceptor is a base ● in a weak system, you would see all four species ● acetic acid→weak, monoprotic acid (acid that donates one proton), most organic acids are weak ● strong/weak convention same as for electrolytes→strong acid=strong electrolytes, weak acids=weak electrolytes November 18, 2015 Bronsted-Lowry Definitions ● Acid→a proton (H+) ● Base→a proton acceptor ● water is not necessary as in Arrhenius ● ex. NaOH (base) + HCl (acid)→HOH + NaCl (water and salt) Acid Base Conjugate Pair ● something that acts as a base on one side of the equation can act as a base on the other side of the equation ● NH₃ + H₂O ⇄ NH₄+ + OH- ● NH₃ is a base on the reactant side, but an acid on the product side, while H₂O is a acid on the reactant side, but a base on the product side ● A conjugate acid has one more hydrogen (one more proton) than the conjugate base, a conjugate base has one less hydrogen (one less proton) than the conjugate acid Types of Acids ● Monoprotic Acid→donates 1 hydrogen, ex. HCl ● Diprotic Acid→donates 2 hydrogens, ex. ● after the addition of .01 mol of HCl to 1L of water the pH decreases to 2 ● after the addition of .01 mole of NaOH to 1L of water the pH rises to 12 Buffer ● a compound that resists a change in pH when a small amount of an acid or bases is added to a solution ● buffers are the shock absorbers of the chemical world, as they help to maintain pH ● they are made by mixing a weak acid and its conjugate base (in the form of salt) ● in the body pH normally 7.35 must stay ±.3 pH units ● acidosis→the pH in the blood approaches 7 (not enough CO2 being removed from the body, due to emphysema, shallow breathing, etc.) as H2CO3 in increased ● alkalosis→pH rises to about 7.4 (too much CO2 removed, can be caused by fever, hyperventilation), caused by a HCO3- overdose ● blood is mostly buffered by H2CO3/HCO3- and also ….. ● [weak acid] = [conjugate base], the pH of the buffer solution is equal to the pKA How do Buffers Work ● suppose a little strong base (OH-(aq)) is added to the blood’s buffer) ● the strong base is going to react with the weak acid, making a weak base ● suppose a little strong acid (H3O+) is added to blood’s buffer) ● the strong acid is going to react with the base, making a weak acid Buffer Capacity ● the extent to which the buffer can resist changes in pH ● the higher the concentration of the weak acid and the conjugate base, the higher the buffer capacity ● the closer the pH of the solution ….. Calculating the pH of a Buffer ● pH = pKA + log [A-]/[HA] ● [A-] is the weak acid and [HA] is the conjugate base ● the one with the more hydrogens is the weak acid November 30, 2015 Chapter 10 Organic Chemistry ● the study of the compounds of carbon ● organic compounds are made up of carbon and only a few other elements ● chief among these are hydrogen, oxygen, and nitrogen ● also present are sulfur, phosphorus, and a halogen (fluorine, chlorine, bromine, or iodine) Organic Structure ● structural formula→shows the atoms present in a molecules as well as the bonds that connects them ● VSEPR model→the most common bond angles are 109.5, 120, 180 ● when carbon is bonded to 4 atoms→109.5, tetrahedral ● when carbon is bonded to 3 atoms→120, trigonal planar ● when carbon is bonded to 2 atoms→180, linear Functional Group ● a part of an organic molecule that undergoes chemical reaction ● they undergo the same types of chemical reactions no matter in which molecule they are found ● to a large measure they determine the chemical ● Alcohols ● primary, secondary, or tertiary→we look at the carbon attached to the OH, and see how many carbons are attached to it Amines ● primary, secondary, or tertiary→how many carbons are directly attached to the nitrogen December 2, 2015 Aldehydes and Ketones ● both contain a C=O (carbonyl group) ● aldehyde →contains a carbonyl group bonded to a hydrogen, in formaldehyde, the simplest aldehyde , the carbonyl group is bonded to two hydrogens ● ketone →contains a carbonyl group bonded to two carbon atoms Carboxylic Acids ● a compound contain a COOH ● carboxyl + hydroxyl group ● condensed formula→CO2H Chapter 11 Condense Formula ● a technique that writes the entire compound on one line ● molecular formula: C5H12→pentane→CH3CH2CH2CH2CH3 ● all are structural isomers of C5H12 Bond Length and Strength ● single are longer than double, and double are longer than triple ● opposite is true with strength Structural Isomers (Constitutional isomers) ● have same formula, but have different structures and properties ● skeletal structures leave out the hydrogen, but we know that carbon always form four bonds ● if you can count the carbons without lifting your finger, it is NOT an isomer ● line angle formulas take into account the geometrical shape, every time you have a bend or end of a line it’s a carbon Classes of Organic Compounds ● alkanes →contain only carbon and hydrogen, and only single bonds ● CnH2n+2 ● n represents the number of carbons Base Names ● KNOW ● use -ane after the base to name an alkane Naming Alkanes 1. find the main chain→start at the end, count the number of carbons connected together 2. circle your main chain 3. are their any side groups? (anything other than hydrogen?) or branches?→have a -yl ending 4. number the main chain to give the lowest number possible to the side group 5. count the number of atoms in the side group, use the right base, add -yl 6. write the carbon it is attached to and separate it from the rest of the chain Know the 8 Most Common Side Groups Read Over Chapters 12 and 13→will not be on exam #3 1. 2-pentene 2. 3-methyl-1-butene 3. 5-chloro-3-octene 1. 3-ethyl- 1, 4-hexadiene 1. 7-bromo-6-chloro-2-dimethyl-5-propyl-3-octyne 2. 7-bromo-6-chloro- double and triple bonds do not allow for free rotation the greater the molecular weight, the greater the boiling point Chapter 12 and 13 Platinum, palladium, nickel and heat are catalysts for alkene to alkane reactions hydration→HO-H