FINAL EXAM REVIEW CHM 2045

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1 FINAL EXAM REVIEW CHM 2045 Disclaimer: Although this review touches on most of the material, it may not cover all the possible material for the exam and should not be used as the only study tool. Go over the exams and quizzes and make sure you understand and know how to solve each one of the problems. Consult Dr. Patino s slides and textbook for further understanding. Also, please keep in mind it is ultimately your responsibility to guarantee that the information here is accurate. *Some material from this review was taken from reviews from previous SI Leaders* YOU CAN DO THIS, I BELIEVE IN YOU. GOOD LUCK!! Matter: anything that has mass and occupies space Atoms = submicroscopic particles, smallest objects that are called matter Molecules = two or more atoms held together by bonds Scientific Approach to Knowledge Hypothesis: tentative interpretation or explanation for an observation Experiment: set of procedure designed to test hypothesis 2 Types of Observations Qualitative Quantitative Law: summary of observations Cannot be violated Allows you to predict future observations Theory: unifying principle that explains facts and laws Q: A chemist mixes sodium with water and witnesses a violent reaction between the metal and water. This is best classified as a) An observation b) A law c) A hypothesis d) A theory

2 Classify each of the following as either gas, liquid or solid. What should you add during each step in this case to get to the next state of matter? What type of reaction would it be? Fixed volume AND shape? Fixed volume with indefinite shape? Indefinite volume AND shape? Classification of Matter Pure substances o = made up of only one element

3 o = made up of two or more elements Mixtures o = uniform throughout o = non-uniform throughout Properties o change = is done without changing composition i.e. changing state of matter (solid, liquid, gas) o change = chemical composition is changed o i.e. burning or bleaching Q: Classify each of the following as physical or chemical change: a) A copper wire is hammered flat b) A nickel dissolves in acid to form a blue-green solution c) Dry ice sublimes without melting d) A match ignites when struck on a flint Q: Three samples, each of a different substance, are weighed and their volume is measured. The results are tabulated below. List the substances in order of decreasing density. Mass Volume Substance I 10.0 g 10.0 ml Substance II 10.0 kg 12.0 L Substance III 12.0 mg 10.0 microliters a) III > II > I b) I > II > III c) III > I > II d) II > I > III Measurement Accuracy: how close to the actual value a measurement gets Precision: how close the repeated measurements are to each other

4 Significant Figures Calculations o Multiplication/Division = lowest number of significant figures Addition/Subtraction = lowest number of significant decimals Q: Perform the calculation to the correct number of significant figures: ( x ) / a) b) 0.12 c) d) Atomic Laws Law of Conservation of Mass Law of Definite Proportions Law of Multiple Proportions Mass is neither created nor destroyed (page 47) All samples of a compound have the same proportion of their constituent elements, regardless of where the samples were taken (page 48) When two elements (A and b) form two different compounds, the masses of element B that combine with 1 gram of element A can be expressed as a ratio of small whole numbers (page 49)

5 Discoveries J.J. Thompson Robert Millikan Ernest Rutherford Measured the charge of the electron Discovered the electron with cathode rays Deduced electrons are negatively charged Developed the nuclear theory of the atom which states an atom is mostly empty space with most of its mass concentrated in the nucleus Atomic Theory (page 50-51) Each element is composed of atoms All atoms of a given element have the same properties that distinguish them from other elements Atoms combine in single, whole-number ratios to form compounds Atoms of one element cannot change into atoms of different elements Isotypes vs. Ions Isotopes: change in # of Ions: change in # of Positive Ions = Negative Ions = Q: Determine the number of protons and neutrons in the isotope Fe-58 a) 26 protons and 58 neutrons b) 32 protons and 26 neutrons c) 26 protons and 32 neutrons d) 58 protons and 58 neutrons Q: A naturally occurring sample of an element contains only two isotopes. The first isotope has a mass of amu and a natural abundance of 60.11% The second isotope has a mas of amu. Find the atomic mass of the element. a) amu b) amu c) amu d) amu

6 Avogadro s # = x 1023 (1 mole) Q: Determine the number of atoms in 1.85 milliliters of mercury. (The density of mercury is 13.5 g/ml) a) 3.02 x 10^27 atoms b) 4.11 x 10^20 atoms c) 7.50 x 10 ^22 atoms d) 1.50 x 10^25 atoms Types of Chemical Bonds Ionic: Covalent: Metallic:

7 Figure 3.10 page 105

8 Q: Name the following: Ba(NO 3 ) 2 : Copper (II) chloride: Cl 2 O: SnO: Manganese (IV) oxide: PCl 5 : Cr(NO 2 ) 2 Calculations/Conversions % Composition: Mass % of element X = (mass of element X in 1 mole of compound) / (mass of 1 mole of compound) 100% Q: List the elements in the compound CF 2 Cl 2 in order of decreasing mass percent composition: a) C > F > Cl b) F > Cl > C c) Cl > C > F d) Cl > F > C Percent Yield = actual yield theoretical yield 100% -Theoretical Yield: calculated amount of product based on chemical equation Actual Yield: actual amount of product produced when doing reaction Limiting Reagent: Excess Reagent: Molarity = moles/liters Dilution = M1V1 = M2V2 Formulas o Empirical: lowest possible number of subscript o Molecular: actual number of atoms in molecule o Structural: lines represent which atoms are bonded to each other Molecular formula = empirical formula x n

9 n = molar mass empirical formula mass Q: The empirical formula of butanedione is C 2 H 3 O, and its molar mass is g/mol. Find its molecular formula Q: Upon combustion, a sample of g composed of only carbon, hydrogen and oxygen produced g of CO2 and g of H2O. What is the empirical formula of the compound?

10 Q: What are the correct coefficients when the chemical equation is balanced? a) 1, 3, 1, 3 b) 1, 2, 1, 1 c) 1, 3, 2, 1 d) 3, 6, 1, 9 PCl 3 (l) + H 2 O (l) H 3 PO 3 (aq) + HCl (aq) Co 2 O 3 (s) + C(s) Co (s) + CO 2 (aq) Solutions Solvent component with highest volume Solute component other than solvent Solution solute + solvent

Electrolytes Dissolve in water Ionic bonds and acids Conduct electricity Strong Electrolytes Dissociate 100% Shown with a one directional arrow Ex: KBr, HCl, NaOH Weak

11 Electrolytes Dissolve in water Ionic bonds and acids Conduct electricity Strong Electrolytes Dissociate 100% Shown with a one directional arrow Ex: KBr, HCl, NaOH Weak Electrolytes Only dissociate slightly Shown with two arrows Less or equal to 5% Ex: HF, NH3 Nonelectrolytes Don t ionize in water Covalent bonds Ex: Sugars Solubility Rules

12 Representations of Chemical Equations Molecular equation -Complete ionic equation Net Ionic Equation Types of Reactions Acid-base Reactions Acid substance that produces H+ Base Substance that produces OH- Neutralization o Acid + Base ---- Salt + Water CH3COOH(aq) + NaOH(aq) CH3COONa(aq) + H2O(l) -Titration Used to quantify the concentration of an unknown acid or base using the opposite with a known concentration Q: Calculate the molar concentration of HCl in a solution if 35.0 milliliters of the acid required 28.9 milliliters of a M KOH solution for titration Redox Reactions Oxidation Is Losing an electron o Reduction Is Gaining an electron o Rules to assigning oxidation numbers Elements in their elemental form have an oxidation number of 0

13 The oxidation number of a monatomic ion is the same as its charge Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions Oxygen using has an oxidation number of -2, except in peroxides, which it then becomes -1 o Ex: H2O vs. H2O2 Hydrogen is -1 when bonded to a metal, and +1 when bonded to a nonmetal o Ex: H2O vs. AlH2 Fluorine always has an oxidation number of 1 The sum of the oxidation numbers in a compound is equivalent to its overall charge o Ex: NO3 - N = +5; O = -2 (+5) + 3 (-2) = -1 Q: What is the oxidation state of carbon in (CO3) 2- a) +4 b) +3 c) -3 d) -2 Q: Determine whether the following reaction is a redox reaction. If it is, identify the oxidizing agent and the reducing agent 2 Mg (s) + O 2 (g) 2MgO (s) MAIN THERMOCHEMISTRY EQUATIONS (MAKE SURE YOU KNOW THESE BY MEMORY) ΔE = Efinal Einitial ΔE = Eproducts Ereactants ΔEsys = Esurroundings ΔE = Efinal Einitial q = C x ΔT q = m x C x ΔT ΔE = q + w w = P ΔT qcal = qrxn ΔH = ΔE + PΔV ΔHrxn = Σ np ΔHf (products) Σ nr ΔHf (reactants)

Calorimetry Coffee-cup o Measures heat of a reaction at a constant pressure q solution = m solution x Cs x T Bomb o Measures internal energy of a reaction at a constant volume q calorimeter = Ccal x

14 Calorimetry Coffee-cup o Measures heat of a reaction at a constant pressure q solution = m solution x Cs x T Bomb o Measures internal energy of a reaction at a constant volume q calorimeter = Ccal x T You need Coffee for the MCAT vs. CATs are the Bomb Q: A chemical system produces 155 kilojoules of heat and does 22 kilojoules of work. What is internal energy of the surroundings? a) 177 KJ b) -177 KJ c) 133 KJ d) -133 KJ Q: How much heat must be absorbed by a 15 g sample of water to raise its temperature form 25.0 Celsius to 55.0 Celsius? (For water Cs = 4.18 J/g*C) a) 1.57 KJ b) 1.88 KJ c) 3.45 KJ d) 107 KJ

15 Hess s Law If chemical equation multiplied by some factor, the H is also multiplied by that factor If chemical equation is reversed, H changes signs If chemical equation can be expressed as a sum of a series of steps, H for overall reaction is the sum of H for each step. Standard enthalpies of formation (use with Lewis Structures) a. H reaction = Σ n Hf (products) Σ n Hf (reactants) MAIN ENERGY EQUATIONS E = hv (MAKE SURE YOU KNOW) v = c/λ (MAKE SURE YOU KNOW) 1 λ = R ( 1 nf 2 1 ni 2) h = J s (Planck s constant) R = m^-1 c=2.998 x 10^8 m/s

16 Electromagnetic Radiation 3.00 x 108 m/s = speed of light Frequency = # of waves that pass a point during certain period of time Greater wavelength = smaller frequency Quantum numbers

17 Q: Which of the following represent valid sets of quantum numbers? For a set that is invalid, explain briefly why it is invalid. a) n = 3, l = 2, ml: -3 b) n = 2, l = 1, ml: +1 c) n = 6, l = 5, ml: 3 d) n = 4, l = 4, ml: +3 Orbital Filling Electrons occupy the lowest energy orbital available (i.e. 1s before 2s) Fill in the following order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s... Pauli exclusion principle: no two electrons in an atom can have the same four quantum numbers Hund s Rule: when filling degenerate orbitals, electrons filled them singly first, with parallel spins Aufbau principle: lower energy levels are filled first Orbitals can only hold 2 electrons each

Electron Configuration Writing all of the energy levels that are filled with electrons (and how many in each) Shorthand = using nearest noble gas preceding the element in

18 Electron Configuration Writing all of the energy levels that are filled with electrons (and how many in each) Shorthand = using nearest noble gas preceding the element in question to shorten the electron configuration Li = [He] 2s1 Periodic trends Zeff: nuclear charge core electrons Increases from left to right Ex: o K: = +1 o C: +6 2 = +4

19 Lattice Energy Lattice energy becomes less exothermic as ionic radius increases Lattice energy becomes more exothermic as magnitude of ionic charge increases Lewis Structures [1] count valence electrons [2] determine central atom [3] distribute bonds between atoms, then add lone pairs to terminal atoms [4] any remaining electrons, after terminal atoms have their octet, go to the central atom Some can have incomplete octets (BF3) or expanded octets (SF6) Bond Lengths Distance between nuclei of bonded atoms: Single bond > double bond > triple bond o Single bond = sigma o Double bond = sigma and pi o Triple bond = sigma and 2 pi The shorter the length, the stronger the bond Polarity Means that one atom has a stronger pull on the electrons than the other atom Differences in electronegativities to determine polarity o Non-polar 0 0.4

20 o Polar o Ionic 2+ Hybridization [1] Count up sigma bonds and lone electron pairs [2] Determine the principal energy level of atom s valence electrons [3] Take s-orbital to hybridize [4] Take as many p-orbitals and d-orbitals (if applicable) needed for same number of orbitals hybridized as sigma bonds/lone pairs

21 State the hybridization and shape for each of the non-hydrogen atoms in the following molecule. Also state the number of sigma and pi bonds. Bond Order Bonding order = (# of electrons in bonding MOs) (# of electrons in antibonding MOs) / 2 Be able to fill in an energy level diagram Gases Pressure: atm is standard unit o 1 atm = 760 mmhg a.k.a. torr o Manometer measures the difference in atmospheric pressure and the pressure of the gas in the vessel Temperature: Kelvin is standard unit o Kelvin = C Density = mass/volume o Can find mass from using PV=nRT to find moles, then convert o STP = standard temperature and pressure o Temperature = 273 K o o Pressure = 1 atm (760 mmhg)

22 Diffusion vs. Effusion Diffusion = collection of molecules spreading from high conc. to low conc. Effusion = collection of molecules escaping through a tiny hole in a vacuum Q: A sample of oxygen was collected by displacement of water at 25 degrees Celsius. The atmospheric pressure was 700 torr. What is the mole fraction of O2? Vapor pressure of H2O is 24 torr at 25 degrees Celsius Q: What are the assumptions/postulates of the Kinetic Molecular Theory?

FINAL EXAM REVIEW CHM IB DR. DIXON

FINAL EXAM REVIEW CHM IB DR. DIXON DISCLAIMER: SOME OF THIS MATERIAL WAS TAKEN FROM OTHER MATERIALS CREATED BY OTHER SI LEADERS. THERE IS A POSSIBILITY THAT THIS REVIEW CONTAINS ERRORS. PLEASE REFER TO YOUR TEXTBOOK, CLASS SLIDES OR YOUR