CHE 101 Principles of Chemistry I: Course Objectives

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1 General Description: CHE 101 Principles of Chemistry I: Course Objectives CHE 101 introduces new concepts and broadens those learned previously; chemical names, formulas, and equations; types of reactions; stoichiometry; thermochemistry; atomic structure and bonding; behavior of gases, liquids, and solids; properties of solutions. Prerequisites: recent high school chemistry or CHE 100 with a grade of C or higher, and MAT 098 with a grade of C or higher. Important Note: Core objectives, listed in bold italics, will be covered in all sections of CHE 101. The remaining objectives listed, as well as an occasional topic within an individual instructor s professional specialty or interest, may be covered at the discretion of the instructor. Unit I. Chemical Foundations 1.1 Distinguish among observations, laws, theories, and hypotheses in the scientific approach to knowledge. 1.2 State and apply rules for taking measurements and significant figures. 1.3 Identify the common SI units for mass (kg), length (m), and time (s) and know the meaning of the following metric prefixes (i.e. memorize and be able to apply): giga, mega, kilo, deci, centi, milli, micro, nano, pico. Students should also memorize and be able to apply 1 cm 3 = 1 ml. 1.4 Perform conversions among different units using dimensional analysis (note: students should be able to apply English to metric and English to English conversion factors, but they do not need to memorize them). 1.5 Interconvert temperatures among Celsius, Fahrenheit and Kelvin scales. 1.6 Solve problems involving density. 1.7 Differentiate between pure substances and mixtures, and correctly use specific terms such as elements, compounds, homogeneous, heterogeneous, states (phases), and solutions. 1.8 Differentiate between chemical and physical properties and changes.

2 1.9 Describe how distillation, filtration, and chromatography can be used to physically separate substances. Unit 2. Atoms, Molecules, and Ions 2.1 State the Law of Conservation of Mass, Law of Definite Proportions, and Law of Multiple Proportions and explain how given data illustrates these laws 2.2 Explain the historical development of the atomic concept leading to the modern view of the atom. This would include being able to describe Dalton s atomic theory, Thomson s cathode-ray tube experiment, the plum pudding model, Milliken s oil drop experiment, and Rutherford s gold foil experiment. 2.3 List the fundamental sub-atomic particles and their basic properties. 2.4 Define isotope, write (and interpret) atomic symbols of the form A Z X, and solve related problems. 2.5 Define molecule, ion, atom, cation, anion, monatomic ion, and polyatomic ion, and solve related problems. 2.6 Describe the type of information conveyed by a molecular formula and a structural formula. 2.7 Correctly use terms associated with the Periodic Table, including Period, Group (Family), representative elements (main group elements), transition metals, lanthanides, actinides, alkali metals, alkaline earth metals, halogens, and noble gases. 2.8 Classify elements in the Periodic Table as metals, nonmetals, or metalloids and list and define the general properties of these classifications. 2.9 Identify the elements which exist as solids, those which exist as liquids, and those which exist as gases at room temperature and pressure Describe the meaning of periodicity as it applies to the Periodic Table and state that the Periodic Table is arranged so that the elements are 1) in order of increasing atomic number, and 2) so that elements in the same column have similar properties Write the name if given the symbol and symbol if given the name of the following elements: 1-56, Pt, Au, Hg, Pb, Rn, Ra, U, and Pu Write the symbol (including the charge) and name for the following monatomic ions: all those of Group 1, Group 2, and Group 17; Al 3+, N 3-, P 3-, As 3-, O 2-, S 2-, Se 2-, Te 2-, Fe 2+, Fe 3+, Cu +, Cu 2+, Ag +, Zn 2+, H +, H -.

3 2.13 Apply the systematic nomenclature (i.e. Roman numeral system) to all possibilities of cations with multiple charges. 2.14Apply the older nomenclature system to Fe 2+ (ferrous), Fe 3+ (ferric), Cu + (cuprous), Cu 2+ (cupric) 2.15 Write the symbol (including charge) and name for the following polyatomic ions: NH 4 +, NO 2 -, NO 3 -, SO 3 2-, SO 4 2-, HSO 4 - (hydrogen sulfate and bisulfate name), OH -, CN -, PO 4 3-, HPO 4 2-, H 2 PO 4 -, CO 3 2-, HCO 3 - (hydrogen carbonate and bicarbonate name), ClO -, ClO 2 -, ClO 3 -, ClO 4 -, C 2 H 3 O 2 -, MnO 4 -, Cr 2 O 7 2-, CrO Write the names and formulas for salts derived from the monatomic and polyatomic ions mentioned above Write the names and formulas for binary molecular compounds: students will expected to apply the following prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca Write the names and formulas for common inorganic acids based upon the anions mentioned above List the seven elements which exist as diatomic molecules. Unit 3. Stoichiometry 3.1 Define atomic mass and solve related problems. 3.2 Solve calculations involving isotopic masses and abundances. 3.3 Solve problems which demonstrate an understanding of the mole concept. 3.4 Calculate the formula mass of a compound. 3.5 Explain and apply the relationship between atomic mass, formula mass, molecular mass, and molar mass. 3.6 Interconvert between the following as it applies to a particular substance: mass, moles, molecules, formula units, atoms, and ions. 3.7 Calculate the percent composition of a compound from its formula. 3.8 Calculate the empirical and molecular formula of a compound from quantitative analytical data.

4 3.9 Differentiate between reactants and products of a chemical equation and recognize symbols for various states of matter (s, l, g, aq) Balance chemical equations by inspection and solve related problems Solve standard stoichiometric calculations Identify the limiting reactant of a chemical reaction and calculate theoretical yields based upon this information Solve problems involving the percentage yield of a reaction. Unit 4. Types of Chemical Reactions and Solution Stoichiometry 4.1 Define the terms solute, solvent, and solution. 4.2 Define the terms strong, weak, and non-electrolyte and solve related problems. 4.3 Define Arrhenius acid and base. 4.4 Define and list common examples of: a. strong acids (HCl, HBr, HI, H 2 SO 4, HNO 3, HClO 4 ) b. weak acids (H 3 PO 4, HF, HC 2 H 3 O 2 [CH 3 COOH]) c. strong bases (Group 1 hydroxides [LiOH, NaOH, KOH, RbOH, CsOH], and larger Group 2 hydroxides [Ca(OH) 2, Sr(OH) 2, Ba(OH) 2 ]) d. weak bases (NH 3 ). 4.5 List and/or recognize all ionic compounds and strong acids as examples of strong electrolytes. 4.6 List and/or recognize weak acids and weak bases as examples of weak electrolytes. 4.7 List and/or recognize the following common examples of nonelectrolytes: sugars (C 6 H 12 O 6, C 12 H 22 O 11 ), alcohols (CH 3 OH, CH 3 CH 2 OH), water. 4.8 State the meaning of molarity. 4.9 Calculate the molarity of a solution (or of the ions in solution for strong electrolytes) and solve related problems Solve molarity problems involving calculating quantities of solute or solution Describe the proper techniques for solution preparation.

5 4.12 Solve problems involving dilutions of solutions Solve stoichiometry problems involving reactants and/or products in solution Recognize and describe the nature of a precipitation reaction, acid-base reaction, and redox reaction Learn and be able to apply the following solubility rules for ionic compounds: Soluble Ionic Compounds: 1. All common compounds of Group 1 ions (Li +, Na +, K + etc ) and the ammonium ion (NH 4 + ) are soluble. 2. All common nitrates (NO 3 - ), acetates (CH 3 COO - ), and most perchlorates (ClO 4 - ) are soluble. 3. All common chlorides (Cl - ), bromides (Br - ), and iodides (I - ) are soluble, except those of Ag +, Pb 2+, Cu +, Hg All common sulfates (SO 4 2- ) are soluble, except those of Ca 2+, Sr 2+, Ba 2+, and Pb 2+. Insoluble Ionic Compounds: 1. All common metal hydroxides are insoluble, except those of Group 1 and the larger members of Group 2 (starting with Ca 2+ ). 2. All common carbonates (CO 3 2- ) and phosphates (PO 4 3- ) are insoluble, except those of Group 1 and NH All common sulfides are insoluble except those of Group 1, Group 2 and NH Predict products and write a balanced molecular equation and net ionic equation for precipitation reactions and acid-base reactions which are double displacement reactions Identify spectator ions in a chemical reaction Define titration and perform calculations related to titrations State and apply the following rules for assigning oxidation numbers: 1. For free elements, the oxidation number is zero. Examples: Na (0), O 2 (0), S 8 (0). 2. For monatomic ions, the oxidation number is the charge on the ion. Examples: Na + (+1), Cl - (-1), Al 3+ (+3). 3. For Group 1A elements, the oxidation number is always +1 when they are in their ionic form. Examples: Na + (+1), Rb + (+1). 4. For Group 2A elements, the oxidation number is always +2 when they are in their ionic form. Examples: Ca 2+ (+2), Ba 2+ (+2).

6 5. For oxygen, the oxidation number is usually a -2. Exceptions: as a free element it is zero; for peroxides it s a -1 (examples of peroxides: Na 2 O 2, H 2 O 2 ); for superoxides it s -1/2 (examples of superoxides: RbO 2, CsO 2 ); with fluorine it s positive (examples of oxyfluorine compounds: OF 2, O 2 F 2 ). 6. For hydrogen, the oxidation number is +1 when combined with other nonmetals and -1 when combined with metals. 7. For fluorine, the oxidation number is always a -1 (except when it is a free element). 8. For binary compounds (those with two different elements) [and in general when two elements are covalently bonded], the element with the greater electronegativity is assigned a negative oxidation number equal to its charge in simple ionic compounds of that element. Example: In PCl 3, Cl is more electronegative than P and since Cl forms a -1 charge in simple ionic compounds (such as NaCl, MgCl 2 ), Cl is assigned an oxidation number of -1 in PCl The sum of the oxidation numbers equals the charge on the molecule or ion Define oxidation, reduction, reducing agent, and oxidizing agent and solve related problems Identify the following in an oxidation-reduction reaction: element getting oxidized, element getting reduced, oxidizing agent, reducing agent. Unit 5. Gases 5.1 Describe what is meant by gas pressure. 5.2 Identify, memorize conversion factors, and know how to interconvert between the following common units of pressure: atm, mm Hg, and torr. 5.3 Identify and, if given the conversion factors, know how to interconvert between the following common units of pressure: Pa, in Hg, and psi. 5.4 State and describe the meaning of Boyle s, Charles s, and Avogadro s Laws. 5.5 Solve qualitative and quantitative problems that involve the application of Boyle s, Charles s, and Avogadro s Laws (Avogadro s law should be understood as stated in Section 5.2 for a gas at constant temperature and pressure, the volume is directly proportional to the number moles of gas and as stated in Section 2.3 at the same temperature and pressure, equal volumes of different gases contain the same number of molecules ). 5.6 Explain the difference between an ideal gas and a real gas.

7 5.7 State the values for standard temperature and pressure and the molar volume of a gas at STP. 5.8 Solve problems involving the ideal gas law (PV = nrt). 5.9 Perform stoichiometric calculations that involve gaseous reactants and/or products State Dalton s Law of Partial Pressures, state the meaning of mole fraction, and solve related problems Describe the difference between effusion and diffusion and solve qualitative problems related to these concepts Use the kinetic-molecular theory to explain the gas laws. Unit 6. Thermochemistry 6.1 Solve problems related to work (in terms of pressure and volume) and the first law of thermodynamics. 6.2 Define exothermic and endothermic reaction and solve related problems. 6.3 Solve problems related to enthalpy. 6.4 Define the terms specific heat capacity (specific heat), heat capacity, and molar heat capacity and solve related problems. 6.5 Solve problems related to calorimetry and describe what a bomb calorimeter is. 6.6 Use Hess s Law to calculate enthalpy changes for a variety of reactions. 6.7 Apply their knowledge of standard states and standard enthalpies of formation to calculate H for a variety of reactions and solve related problems. Unit 7. Atomic Structure 7.1 List the eight regions of the electromagnetic spectrum in the designated order and perform calculations involving frequency, wavelength, and energy (note: students should memorize E = hν and c = νλ and be able to apply Planck s constant and the speed of light if given)

8 7.2 Define photon and relate the photon to the dual nature of electromagnetic radiation. 7.3 Describe the relation between electronic transitions and line spectra. 7.3a Solve quantitative problems related to electronic transitions in the hydrogen atom. 7.4 Explain in a qualitative way the Bohr model of the hydrogen atom. 7.5 List the four quantum numbers and relate them to electronic structure by using them to predict relevant information about shells, subshells, orbitals, and electrons. 7.6 Differentiate between a Bohr orbit and a quantum mechanical orbital. 7.7 Sketch any s, p, or d orbital. 7.8 Write electronic configurations (i.e. Mg = 1s 2 2s 2 2p 6 3s 2 ) for all elements which strictly follow the Aufbau principle, Pauli exclusion principle, and Hund's rule and solve related problems. 7.9 Write electron configurations for the anomalies Cr, Mo, Cu, Ag, Au and write electron configurations for all other anomalies in which adequate hints are provided and solve related problems Use the Periodic Table to write noble gas abbreviated electron configurations (i.e. Mg = [Ne]3s 2 ) Draw orbital diagrams based on electron configurations (i.e. Mg = 1s 2s 2p 3s and use them to predict paramagnetism and diamagnetism. ) 7.12 Define paramagnetic and diamagnetic Explain the concept of effective nuclear charge and the trend which exists across a period Define ionization energy and electron affinity and be able to write equations related to each Use the Periodic Table to predict trends in ionization energies, electron affinities, and atomic sizes and to provide adequate explanations for these trends. Unit 8. Bonding 8.1 Determine the number of valence electrons for main group elements.

9 8.2 Draw Lewis dot symbols for atoms and ions of main group elements. 8.3 Describe the difference between ionic and covalent bonding modes. 8.4 Recognize that bonding is an example of a model. 8.5 Predict which combinations of elements tend to bond ionically, polar covalently, and nonpolar covalently. 8.6 State and explain the effect of a loss or gain of electrons on atomic size. 8.7 Describe the meaning of lattice energy and solve qualitative problems related to the concept. 8.8 Write the electron configurations of ions. 8.9 Define electronegativity and apply it towards determining bond polarity Determine the formal charge on any atom within a molecule or ion Draw Lewis structures for molecules and polyatomic ions and to recognize when to apply the octet rule model and when to apply the model which satisfies formal charge rules Draw resonance structures for molecules and ions and solve problems related to the resonance concept Use the Valence Shell Electron Pair Repulsion Theory (VSEPR) to predict electron pair geometries for the following electron pair arrangements: linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral Use VSEPR to predict bond angles and molecular geometries, and to sketch shapes of molecules using perspective drawings Predict whether molecules are polar (have a dipole moment) or nonpolar (do not have a dipole moment) Use the valence bond theory (localized electron model) to predict hybridization and solve related problems Describe the difference between sigma and pi bonds and be able to sketch pi bonds State the relative bond lengths of similar single, double, and triple bonds and solve related problems.

10 8.19 Describe basic molecular orbital theory, in particular how it contrasts with valence bond theory Define and explain the difference between bonding and antibonding orbitals Draw MO diagrams for period 1 and 2 homonuclear diatomics and use them to evaluate bond order. Unit 9. Liquids and Solids 9.1 List the various intermolecular attractions in liquids and solids (dipole-dipole, London dispersion, hydrogen bond), describe the nature and relative strength of each, and identify which compounds exhibit which intermolecular attraction. 9.2 Predict and explain how intermolecular attractions affect the physical properties of liquids and solids. 9.3 List the characteristics of molecular, network covalent, ionic and metallic solids. 9.4 Describe the relationship among temperature, vapor pressure, and boiling point. 9.5 Define heat of vaporization and heat of fusion and solve related problems. 9.6 Differentiate between the terms melting, freezing, boiling, evaporation, condensation, sublimation, and deposition. 9.7 Draw and interpret a heating/cooling curve for a substance. 9.8 Describe the nature of a supercooled liquid and a superheated liquid. 9.9 Solve quantitative problems related to heat of fusion and heat of vaporization Define normal melting point and normal boiling point Interpret phase diagrams Define the terms critical point and triple point as it applies to phase diagrams. Unit 10. Properties of Solutions 10.1 Differentiate between the following ways of expressing the concentration of a solution: molarity, molality, mole fraction, and mass percent.

11 10.2 Solve quantitative problems related to molarity, molality, mole fraction, and mass percent Predict the solubility of various solutes in various solvents based on polarity and intermolecular forces Define the terms saturated, unsaturated, and supersaturated Predict the effects of changes in temperature and pressure on the solubility of solids and gases and solve related problems Solve both qualitative and quantitative problems related to the colligative properties of freezing point depression and boiling point elevation (quantitative problems for nonelectrolytes only).