Study Guide for Exam 1
Study Guide for Exam 1 CHEM 111
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This 15 page Study Guide was uploaded by Julia Notetaker on Friday October 7, 2016. The Study Guide belongs to CHEM 111 at University of Massachusetts taught by David Adams in Fall. Since its upload, it has received 5 views. For similar materials see /class/232341/chem-111-university-of-massachusetts in Chemistry at University of Massachusetts.
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Date Created: 10/07/16
CHEM 111 STUDY GUIDE- EXAM 1 CHAPTER 1: ATOMIC SCALE • Properties of matter depend on the types M(What we c and arrangement of observe) molecules on atomic scale o Ex) Pop a red balloon Levels of without heat- Matter normal burst Symbolic Particulate o Pop a red liquid to imagine) balloon with water gas) heat- louder burst and flames o Heat (flame) = reactant that starts the reaction Definitions: • Nomenclature: a set or system of names • Stoichiometry: measures quantitative relationships and is used to determine the amount of products/reactants that are produced/needed in a given reaction • Physical Change: changes affecting the form of a chemical substance but not its chemical composition • Chemical Change: any change that results in the formation of a new chemical substance • Energy: power derived from the utilization of physical or chemical resources Classifying Matter • Definitions: Matter o Elements: § Basic unit of matter (materials) § 1 type of atom § Cannot be broken down Substances Mixtures § Things on periodic table o Periodic Table: § Total amount of elements: 118 § 90 % earths crust: Elements Compounds Physical Process O2, Si, Al, Fe, Ca § 90% body mass: C, H, O Chemical o Atoms: Process § Smallest chemical unit of matter or indivisible unit of an element o Molecules: § Formed when 2+ atoms join together chemically o Compounds: § At least 2+ atoms of different elements bonded together in defined ratio • Can be broken down into smaller units via chemical reactions § When gong from molecules à compound: properties of the element change KEY NOTE: All compounds are molecules but not all molecules are compounds States of Matter Solid Liquid Gas • When moving from one state of matter to another: o NOT changing chemical composition o Changing: state of energy § Done by either addition or removal of energy § Known as phases § Considered physical changes Physical State Packing Interaction Kinetic (motion) Between Energy of Molecules molecules Solid Close-Packed Strong Least Liquid Intermediate Intermediate Intermediate Gas Loosely- Packed Weak Greatest • Definitions: o Pure Substances: § Collection of matter that has same chemical composition throughout § Can consists of single element OR compound § Have unique physical properties § CANNOT be separated by physical techniques § Ex) Iron (Fe), table salt (NaCl) o Chemical Change: § Involves change in the chemical composition of the material § Ex) Iron turning to rust in the presence of air • Combustion of Hydrogen (H ): H2 + O2 = 22 0 2 Physical Properties/Process Chemical Properties/Process Density Reacting with other compounds/molecules Color Combustion/burning of air Melting Point Boiling Point Condensation Evaporation Dissolving Substances à Mixtures = NO chemical change Elements à Compounds= Chemically Changing Properties Definition: • Mixture: 2+ substances in variable amounts that have NOT chemically reacted o Ex) NaCl and H O 2 Mixtures Homogenous Heterogenous Uniform Constant Composition NON-Uniform Consists of multiple Consists of single aka) solutions phases phase Ex) Oil and water, sand and water Ex) sugar and water, salt and water • Separating Mixtures: Decantation Density (moving mixture Centrifugation frto another) er Boiling Point Distilation State of Matter Filtration Intermolecular Chromatography Forces Vapor Pressure Evaporation Magnetism Magnets Solubility Filtration KEY NOTE: When chemically reacted CANNOT use physical means of separation Temperature K = °C + 273.15 • Allows for the transition from °C àK or K à °C • Celsius and Kelvin share a direct relationship Density D = M (g) / V (mL) • Allows for the conversion between mass and volume of a substance • Matter with high density will sink • Matter with low density will float • Volume does NOT affect the density • Composition CAN affect the density • Classroom example below o Regular Distilled Water: § Diet Coke floats because it contains less sugar = less mass § Diet coke density < density than distilled water = floats § Regular Coke and Orange Soda > density than distilled water= sinks o Salt Water: § Coke, Orange Soda, and Diet Coke density < density of salt water = floats § Example of composition affecting density Diet Coke Diet Floats Coke Floats Tub of OSoda RCoke r Distilled Floats Floats Water Regular Orange Coke Soda TSalt Sinks Sinks Water KEY NOTE: Ice is less dense than water, which is why we have ice burgs Scientific Notation • Used to represent very large numbers x o General Form: N x 10 o N = number between 1 and 10 o X = (+) or (-) integer § Depends on whether number is greater than 1 (positive) or less than 1 (negative) • Ex) 602200000 = 6.022 x 10 8 o Positive integer because move decimal point 8 places to the left -8 • Ex) .0000000625 = 6.022 x 10 o Negative integer because move decimal point 8 places to the right Scientific Figures Definitions • Significant Figure: digits within measurement (certainty and uncertainty) o Certainty in measurements is limited by the instrument used to make the measurement • Sig-Fig Rules o 1. Zeroes between non-zero digits = significant § Ex) 1005 = 4 sig-figs o 2. Zeroes at the beginning of a number = NOT significant § Ex) 0.02 = 1 sig-fig (#2) § Ex) 0.0026= 2 sig-figs (#2, #6) o 3. Zeroes at the end of a number are significant IF the number has a decimal point § Ex) .0200 = 3 sig-figs (#2, #0, #0) § Ex) 3.0 = 2 sig-figs (#3, #0) • Rounding Sig-Figs o Round up when number is </= to 5 • Adding/ Subtracting Sig-Figs o Final answer should have same number of decimal places as value in data set with the lowest number of sig-figs • Multiplying/Dividing Sig-Figs o Final answer should have the same number of sig-figs as value in data set with the lowest number of sig-figs Precision and Accuracy Definitions: § Precision: how close a set of measurements are to each other § Accuracy: how close measurements are to real value CHAPTER 2: ELEMENTS AND COMPOUNDS KEY NOTES: Properties of matter depend on the type of bonds Definitions: § Atoms o Smallest chemical unit of matter § O § H § N o Consists of 3 subatomic particles § Protons § Neutrons § Electrons o Uncharged atoms have nucleus surrounded by electrons § Atomic Composition o Protons: (+) – inside nucleus o Neutrons: inside nucleus o Electrons: (-) – outside nucleus § Element Symbol o Single letter: represents the element/compound § Ex) Carbon = C § Atomic Number (Z) o Number of protons o All atoms of the same elements will have the same number of protons § Atomic Mass (A) o Number of protons + neutrons o Can change for different atoms of same element § Isotopes o Atoms with the same atomic number (protons) BUT different atomic masses (neutrons) § Total percentage abundance of all isotope of elements should be 100% § Atomic Weight o Mass collection of atoms of an element has average value § Average Mass = Atomic Weight o Atomic Weight= ((% abundance isotope A/ 100) + Mass isotope A) + ((% abundance isotope B/100) + Mass isotope B)) § Periodic Table o Organized according to their properties o Classifies metals and non-metals o 114 known elements § 83 stable and in nature § 7 in nature but not radioactive § 24 not on earth (made in labs) o Organized by § Groups • Total of 18 § Rows • Total of 7 • Row 57 and row 89- moved to separate section to allow periodic table to fit on one piece of paper) § Compounds o 2+ atoms of different elements bonded together in defined ratio § Physical and chemical characteristics of the compound from the properties of the individual elements that contribute the compound § Valence Electrons o All elements have certain number of total electrons § Equal to the number of protons in neutral atoms o 2 categories § Valence electrons: outermost layer/shell of the element • When shells get farther away from the nucleus the energy increases § Core electrons: any other electrons in the atom § Ionic Bond o Bond between anion and cation Ionic Compounds § Cations (+) o Positively charged ions o Atom loses an electron § Anion (-) o Negatively charged ions o Atom gains an electron o Ex) NaCl § Na (sodium) has 1 valence electron in outermost shell § Cl (chloride) has 7 valence electrons in outermost shell § Easier to lose 1 than to give 7 § Na give chloride 1 electron causes chloride to have 8 valence electrons in outermost shell = feels “noble” and is complete § Na gives 1 electron = cation § Cl gains an electron = anion § Predicting atom’s charge o Atoms want to feel “noble” (noble gases are in group 8A of periodic table) o Noble means that they have filled valence electron shells (8 max) § Refer to previous NaCl example o Known as the Octect Rule KEY NOTE: In general: • Metals: lose electrons = cations • Non-metals: gain electrons = anions Definitions: • Monoatomic Ions: made of 1 atom o Formed when an atom gains or loses 1 electron o When naming: name of element goes first followed by ion /cation § Ex) • K+: potassium ion/cation • Al 3+: aluminum ion/cation § Always located in the (A) groups- elements always lose/gain same number of electrons each time o Transition Metal Cations: name of element followed by the roman numeral then cation/ion § Ex) • Cr + : chromium (II) cation 3 • Co + : cobalt (III) cation § Located in the (B) groups- number of electrons lost/gained can change Naming Monoatomic Ions • Anion: root name of element followed by “ide” MEMORIZE o F- : Fluoride o Cl- : Chloride o Br- : Bromide o I- : Iodide o O -: Oxide 2 o S : Sulfide o H-: Hydride Naming Polyatomic Ions • MEMORIZE: Negative Charge Positive Charge 2- + o CO : 3arbonate - NH 4: Amonium o AH Co3 : Acetate o NO : Nitrate - o MnO : Pe4manganate o SO : 4ulphate o Cr O 2 7ichromate o PO : Phosphate 4 o ClO : 4erchlorate o CrO : 4hromate o ClO : Chlorate 3 IONIC COMPOUNDS: FORMULA & NAMING • Ionic Compounds are represented in simplest ratio o Ratio= Cation : Anion o Cation written first followed by anion o In neutral ionic compounds the cation charges = anion charges § Ex) Neutral Ionic Compound: NaCl • Cation: Na (+1) • Anion: Cl (-1) • Ratio = 1:1 § Ex) Write formula for Barium Chloride • Cation: Barium (Ba), (+2) o (+2) is due to positioning on periodic table. (Ba) is located in group 2A § 2A = 2 electrons in outermost shell § Easier to lose 2 electrons than gain 6 electrons (to make 8 which equals complete outer electron shell) § Loses 2 electrons = Cation • Anion: Chloride (Cl), (-1) o (-1) is due to positioning on periodic table. (Cl) is located in group 7A § 7A = 67electrons in outermost shell § Gains 1 electrons versus losing 7 electrons § Gain electrons = anion • Charges are currently unequal o Ba (+2) o Cl (-1) o Need to make balanced § Cl (-1) x 2 = Cl (-2) àbalances equation • BaCl2 à final answer § Ex) Write the name for: FeCl 3 • Cation: Fe = Iron o (2+ or 3+) o Since we know (3) Chlorides were used- choose (Fe +3) • Anion: Cl = Chloride o (-1) • Balance: o Fe: (+3) o Cl: (-1), (-1), (-1) • Any non-minerals? o Fe is within B blocks so used roman numerals • Iron (III) Chloride à final answer Covalent/Molecular Compounds Definitions: • Covalent Bond: when two or more non-metals bond with each other (sharing of electrons) • Ionic Bond: Held together by (+) and (-) charges- metals + non-metal • Covalent Bond: Held by sharing of electrons- non-metals only • Network Covalent Bond: atoms held together by a network of covalent bonds • Represent Covalent Bonds: o Molecular Formula: H 0: s2ows how many electrons each element has o Structural Formula: H-O-H: shows how electrons bonds are held together o Ball- and- stick Model: shows geometry of bond o Space-filling- Model: shows sizes of molecules Naming Compounds • Rule 1: first word = first element in compound • Rule 2: second word = second element in compound + ‘ide” • Rule 3: prefixes show how many atoms of each type there are • Rule 4: DON’T use “mono-“ o Ex) BCl 3 § B: Boron § Cl: Chloride § (3) = “tri” § Boron Trichloride o Ex) Dinitrogen trioxide § “Di”= 2 § Nitrogen § “Tri” = 3 § Oxide = Oxygen § N 2 3 Allotropes Definition: • Allotrops: different forms of the same element o Carbon can be network bonded to make diamonds and graphite § Both look physically and structurally very different but both use all (C) atoms § Each carbon is covalently bonded to 4 other Carbons Inorganic Acids • Produce H+ ions when dissolved in water KEY NOTE: In general, (H) is the first element in the formula of inorganic acids o Ex) HCl: hydrochloric acid o Ex) HBr: hydrobromic acid • Use “ous” when oxygen atoms > than hydrogen atoms o Ex) H SO2: 4ulfric Acid CHAPTER 3: STOICHIOMETRY Definition: • Stoichiometry: relationship between quantities of materials in a chemical reaction 23 • Avogadro’s Number: 1 mole = 6.022 x 10 atoms o 1 mole = exact number of atoms present in a sample that weighs the elements o Ex) Zinc has atomic mass of 53.39 (g) à 1 mole of zinc weighs 53.39 (g) • Molar Mass: Mass of 1 mole of atoms o 1 mole of an element has mass (g) = to atomic weight • Molar Mass of a molecule: sum of the molar mass of all of the atoms in the molecule o Molar mass of a compound = molecular weight or formula weight of the compound o Ex) Molar Mass of NaCl § Atomic weight of Na= 22.90 g/mole § Atomic weight of Cl= 35.453 g/mole § Molar mass = 22.90 + 35.453 = 58.443 g/mole § 1 mole of NaCl weighs 58.443 o If element has coefficient in from of element symbol then multiply atomic weight by coefficient and then add with other elements • Percent Composition by Mass: percent by mass of each element in the compound o Aka: percentage by mass o Percentages should equal 100 when added o % Composition of an element = ((number of atoms in the element) x (molar mass of the element) / molar mass of the compound) x 100 Empirical Rule and Molecular Formula Definition: • Empirical Formula: simplest whole number ratio of elements in a compound o DOES NOT provide information of number of atoms in molecule o Molecular formula DOES provide information of number of atoms in molecule o Ex) Glucose = C H6O 12 6 § Divide all subscripts by 2 § CH 2à simplest whole number ratio of elements in compound • Finding Empirical Formula from % Mass o Pure compound always consists of same element combined in same proportion by weight o Ex) Hydrazine = 87.42% nitrogen, 12.58% hydren § M Hydrazine 2.06 g/mole § 1. Determine empirical formula • Convert (% to grams) o Nitrogen = 87.42 g o Hydren = 12.58 g § 2. Convert (g) to moles by dividing atomic mass of element • Nitrogen: 87.4 g / 14.01 g = 6.240 moles • Hydren: 12.58 g / 1.01 g = 12.5 moles § 3. Divide each answer from step 2 by the smallest answer in step 2 • 6.240/6.240= 1 Nitrogen • 12.5/6.240= 2 Hydren § 4. Resulting moles = subscripts in empirical rule § NH 2(don’t write 1 because understood its 1 without it) • Finding Molecular Formula from Empirical Rule o Ex) M Hydrazine 2.06 g/mole § Step 1: calculate molar mass of empirical rule • NH 2 empirical rule, M Nitrogen 14.01 g/mole, M Hydrogen= 1.01 g/mole • 1(14.01) + 2(1.01)= 16.03 g/mole § Step 2: Divide molar masses that’s given by answer from step 1 • M Hydrazine .06 g/mole) / M NH2 (16.03 g/mole)= 2 § Step 3: Multiple subscript for each element in empirical formula by whole number from step 2 • NH 2 = 2H 4 molecular formula for hydrazine Introduction to Balancing Chemical Equations Definitions: § Balanced: same number and kinds of atoms on each side of reaction § Reactants: left side of the reaction § Produces: right side of the reaction § Stoichiometric Coefficient: numbers in front of atoms § Physical State: o (g): gas o (s): solid o (l): liquid o (aq): aqueous à dissolved in water § Law of conservation of matter: matter can neither be created nor destroyed. o Atoms conserved in chemical reaction § Balancing Equation o Ex) C3H 8(g) + 0 2(g) à CO 2g) + H 02 § Step 1: look at how many of each element you have • C = 3 à C = 1 • H = 8 à H= 2 • O = 2 à 0=3 • NOT balanced § Step 2: start by balancing least number on reactant side (DO NOT choose Oxygen because oxygen appears once on reactant and twice on product- easier to save for last) • Balance C: o C H3 (8) + O 2g) à 3CO (g2 + H 0 2 § Step 3: balance next element • Balance H: o C H3 (8) + O 2g) à 3CO (g2 + 4H 0 2 § Step 4: balance next element • Balance O: o C H3 (8) + 5O (2) à 3CO (g)2+ 4H 0 2 § Step 5: Check to see if completely balanced o C H3 (8) + 5O (2) à 3CO (g)2+ 4H 0 2 • C = 3 à C = 3 • H = 8 à H= 8 • O = 10 à 0=10 • Balanced Reaction Stoichiometry Definitions § Reaction Stoichiometry: study of the relationship between the amount of reactions and the products on a macroscopic scale o Ex) Determine the amount of H SO (2ole4) consumed and the amount of MgSO (m4 es) produced when 3.5 moles of Mg N reacts 3 2 § Step 1: Balance equation: Mg N 3 H2 O à2MgS4 (NH ) 4 4 2 4 • Mg N + 4H SO à 3MgSO + (NH ) SO 3 2 2 4 4 4 2 4 § Step 2: Multiple each coefficient for compound involved in (consumption, produced, and reaction) by 3.5- process of converting to moles • 1 x 3.5 = 3.5 moles of Mg N 3 2 • 4 x 3.5 = 14 moles of H S2 4 • 3 x 3.5 = 11 moles of MgSO 4 • Therefore 14 moles of H SO 2re 4 nsumed and 11 moles of MgSO ar4 produced Limiting Reagent/Reactant: Definitions: • Reactant: one that runs out first in chemical reactions o When the reactants run out the chemical reaction stops • Excess reactant: what’s left over once the reaction stops o Ex) In a complete combustion reaction of C H a 14 m10e of C H is 4 10 combined with 120 moles of oxygen. What is limiting reactant? § Atomic Weight (C) = 12.01 g/mole § Atomic Weight (H) = 1.01 g/mole § Step 1: balance equation: C H4+ 10 à C2 + H 02 2 2C 4 +103 O à 2 CO + 102H 0 2 • Step 2: convert from grams to moles (if needed)- doesn’t apply in this situation • Step 3: determine which is limiting reactant o 12 moles of C H x4(110 oles of O / 2 mo2es of C H ) = 784 10 moles of O 2 • Therefore, 12 moles of C H 4r10uire 78 moles of O . We h2ve 120 moles of O so C H is the limiting reactant. 2 4 10 Percent Yield Definition • Percent Yield: efficiency of a reaction is given by % yield o Cannot exceed 100 • Experimental Yield: actual amount of product produced in the experiment • Theoretical Yield: calculated amount of product from a given amount of reactant o Theoretical yield will NEVER exceed experimental yield • % Yield = (experimental yield/ theoretical yield) x 100 • Ex) consider the reaction of Pb(NO ) 3 2h NaCl to form PBCl and N2No . 3 If 24.3 g of Pb(NO 3 2reacted with excess NaCl and 17.3 g of PBCl is 2 isolated. What is the theoretical yield? • M Pb(NO3)2: 331. 2 g/mol • M PBCl2: 78.1 g/mol o Step 1: balance equation: balanced § Pb(NO ) 3 2 + NaCl à PBCl + 22 aNo 3 o Step 2: theoretical yield of PBCl 2 § (24.2 g/mole of Pb(NO ) /1)3 2 1 mole of Pb(NO ) / 331.23 2 g/mol) x (1 mole PBCl /1 m2le of Pb(NO ) ) x 2783 2g of PBCl /21 mol) = 20.3 g of PBCl 2 o Step 3: % yield § % yield = (17.3)/(20.3) x 100 = 85.2% yield Hydrated Compounds Definition § Hydrated Compound: Ionic- Compound that has a well-defined amount of H 0 2 trapped in crystalline solid § Water of Hydration: water associated with the compound o Ex) 32.86 g sample of hydrate of CoCl was heat2 thoroughly in a porcelain crucible unit its weight was constant. After heating, 17.93 g of the dehydrated compound remained in the crucible. What is formula of hydrate? § Step 1: • Wt. of CoCl =2 29.84 g/mole • Wt. of H 2= 18.02 g/mole • CoCl 2 . 0 à2CoCl + H 02 2 § Step 2: mass of water lost • 32.86- 17.93 = 14.93 g of H 0 lo2t § Step3: Mass of water from (g) to moles • (14.93 g of H 02/1 ) x (1 mole of H 0 / 28.02 g/mol) = 0.8285 mol § Step 4: Mass of CoCl 2 • (17.93 g of CoCl /12 x (1 mole of CoCl /129.2 g) = 0.1381 mol § Step 5: ratio of moles of CoCl to2 0 2 • (0.8285/0.1381)= 6/1 = CoCl to 6 H2 2
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