Chem 110 final PDF

Preview

Summary

Review notes for final exam...

Description

CHEM 110B Learning Objectives

Skills Check 1. Perform calculations with measured values and express the resulting value with the appropriate number of significant figures. 2. Convert between metric units. Lecture 2 1. Define the term isotope and interpret atomic symbol notation. 2. Calculate the numbers of protons, neutrons, and electrons for different isotopes and ions. 3. Calculate the average atomic mass for an element, given its isotopes and percent abundances. 4. Describe the basic features of mass spectroscopy (MS) and interpret simple MS data. This is done in the spirit of seeing experimental data and relating to biological applications. 5. Using knowledge of Avogadro’s Number and the mass-mole relationship, calculate masses of individual atoms or molecules, mole quantities from mass amounts and vice versa, and particle quantities from mass or mole amounts. Lecture 3 1. Identify the “system” and “surroundings” in a given chemical/physical process. 2. Apply the First Law of Thermodynamics to calculate the change in internal energy of system, the heat change of a system, or the change in work of a system given the other two. 3. Assign positive and negative signs to internal energy, heat, and work for specific processes. 4. Define enthalpy, differentiate endothermic and exothermic process, and apply those terms appropriately to a system. 5. Define state function and give specific examples. 6. Use Coulomb's Law to rank the electrostatic potential energies associated with interacting charges based on their charge magnitudes and distance from each other. Lecture 4 1. Define wavelength, frequency, and photon energy, and explore the relationship between wavelength and frequency, between energy and wavelength, and between energy and frequency. 2. Rank the relative order of wavelength (frequency, photon energy) for different regions in electromagnetic spectrum. 3. Rank the relative order of wavelengths for the colors in the visible spectrum. 4. Calculate the frequency or wavelength of a photon, the energy of one or multiple photons, and the total number of photons in a pulse of light of a given energy and wavelength. 5. Connect line spectra to the energy/wavelength of absorbed/emitted photons. 6. Define the energy of an electron (E given off when an electron is put into an orbit). 7. Use the energy level diagram of the hydrogen atom to compare the relative wavelength/frequency/energy for different electronic transitions. 8. Identify the “n” values that are associated with ground state and ionization for a hydrogen atom in the Bohr Model. 1

9. Interpret absorption spectra to determine the wavelengths that are most/least absorbed by a substance. Lecture 5 1. Describe the meaning of orbitals and be able to identify orbitals based on their shapes. 2. Determine whether a set of quantum numbers is an allowed combination, and if so, be able to identify the shell and subshell of the orbital defined by that combination. 3. Relate quantum numbers to the size, shape, and orientation of an orbital. 4. Define nodes and identify the number of nodes in an orbital. 5. Describe electron density in regions of space using electron density plots 6. Define the term degenerate and recognize whether sets of orbitals are degenerate. Lecture 6 1. Describe the size, shape, and orientation of orbitals and the spin of the electron given a set of quantum numbers. 2. Describe the Pauli Exclusion Principle and apply it to the electron configurations of atoms. 3. Explain the differences in orbital energy diagrams for the H atom vs. many-electron atoms. 4. Define effective nuclear charge, Zeff, and explain why it differs from the actual nuclear charge, Z. 5. Apply the Aufbau Principle and Hund’s rule to write the ground state electron configuration of any main group element and the first row transition metal elements. 6. Differentiate excited state and ground state electron configurations 7. Relate between electron configurations and the quantum numbers that describe those electrons. Lecture 7 1. Rank neutrally charged atoms in terms of atomic radius. 2. Describe how periodic trends in Zeff influence atomic radius and ionization energy. 3. Differentiate between ionization energy and electron affinity and write the chemical processes associated with these energies. 4. Interpret a plot or a series of successive ionization energies and determine the number of valence electrons in the atom. 5. Compare the ionization energies of different elements based on their positions in the periodic table. 6. Broadly describe the differences in ionization energies of the metals, nonmetals, and noble gases. 7. Predict the charge of an element’s most common ion using the periodic table. 8. Compare the sizes of ions to their corresponding neutrally charged atoms. 9. Rank ions within an isoelectronic series in terms of ionic radius.

2

Lecture 8 1. Differentiate and relate between molecular formula and empirical formula. 2. Using knowledge of Avogadro’s number and the mass-mole relationship, calculate masses of individual atoms or molecules, mole quantities from mass amounts, mass quantities from mole amounts, and particle quantities from mass or mole amounts. 3. Calculate molecular (formula) weight and determine mass percent of an element in a compound. 4. Given the results of a percent composition of a substance, determine the empirical formula for the compound. If sufficient additional information is provided, such as the molar mass, determine the molecular formula as well. 5. Determine the molar mass of a compound from its mass spectrum. Lecture 9 1. 2. 3. 4.

Recognize and describe the types of bonding. Identify whether a compound is ionic or molecular. Describe the octet rule and draw Lewis dot symbols for elements/ions. Define lattice energies and use Coulomb’s Law to understand lattice energy trends in ionic solids. 5. Predict the number of covalent bonds an element will typically form (i.e., its valence). 6. Correlate the bond order, bond length, and bond strength of different covalent compounds. Lecture 10 1. Use electronegativity trends across the periodic table to compare the electronegativity values of elements and determine whether bonds are polar. 2. Rank bonds based on polarity. 3. Apply the octet rule and rules of valence to draw Lewis structures. Lecture 11 1. Assign formal charges to atoms in molecules and label formal charges in Lewis structures. 2. Select the best Lewis structure from multiple possibilities based on formal charges. Lecture 12 1. 2. 3. 4.

Recognize molecular structures in which resonance will exist. Draw a molecule’s different resonance structures. Identify molecules that may violate the octet rule. Draw Lewis structures for hypervalent molecules, molecules with incomplete octets, and molecules with odd numbers of electrons.

3

Lecture 13 1. Draw and interpret carbon backbone (i.e., skeletal or line-angle) representations of organic molecules. 2. Apply the electron pair repulsion model (VSEPR) and use Lewis structures to determine the electron domain geometries and molecular geometries of atoms in a molecule. 3. Predict the effects of lone pairs and multiple bonds on bond angles. 4. Analyze the shape and bond angles of each central atom in a large molecule. Lecture 14 1. Use molecular geometry and electronegativity differences to determine whether a molecule will have an overall dipole moment. Lecture 15 1. Determine the hybrid orbitals involved in covalent bonding using Lewis structures, VSEPR shapes, and with consideration for resonance. 2. Differentiate between sigma and pi bonds, and recognize the types of orbitals involved in these bonds. Lecture 16 1. 2. 3. 4.

Identify orbitals used to form multiple bonds. Explain the restricted rotation in a multiple bond. Count the number of sigma and pi bonds in a molecule. Predict when pi bonds and lone pairs will be localized and when they will be delocalized.

Lecture 17 1. 2. 3. 4. 5.

Classify a molecule as hydrocarbon (alkane, alkene, alkyne, or aromatic). Identify constitutional isomers. Identify geometric isomers and label them as being cis or trans. Explain how melting point varies for different geometric isomers. Identify functional groups in a molecule.

4

Lecture 18 1. Identify functional groups in molecules (including alcohols, ethers, amines, and carbonylcontaining groups such as carboxylic acids, esters, amides, ketones, and aldehydes). 2. Describe the basic structure of an amino acid and be able to locate the alpha carbon, carboxylic acid group, amine group, and side chain. 3. Use the structures of amino acid side chains to describe the properties of each amino acid and classify the side chain as being polar or nonpolar. 4. Recognize chiral carbon atoms (i.e., stereocenters) and explain how stereoisomers are structurally distinct from one another. 5. Describe the structural features of the peptide bonds that connect amino acids together in proteins. Lecture 19 1. Differentiate between intramolecular and intermolecular forces. 2. Explain the differences between solids, liquids, and gases at the molecular level, and relate kinetic energy and temperature to IMF strength and phases of matter. 3. Use melting point and boiling point data to determine relative IMF strength. 4. Describe and identify ion-ion, ion-dipole, ion-induced dipole, dipole-dipole, dipole-induced dipole, and dispersion interactions. 5. Explain how induced dipoles allow nonpolar molecules to participate in intermolecular interactions. 6. Define polarizability and describe how molecular size and shape influence dispersion forces. Lecture 20 1. 2. 3. 4.

Identify the types of IMFs present between any chemical species. Identify whether molecules can form hydrogen bonds. Identify and label hydrogen bond donors and acceptors. Predict trends in IMF strength for sets of molecules based on size, shape, polarity, and ability to form hydrogen bonds. Also be able to predict trends in melting/boiling points.

Lecture 21 1. Describe the chemical structure of a polymer and be able to identify the monomers from which it is constructed. 2. Describe the chemical structure and properties of the peptide bond; understand how resonance causes the peptide bond to be flat and especially polar. 3. Identify the N- and C-termini of a polypeptide chain, even when the termini are not explicitly shown. 4. Identify the bonds in the peptide backbone that are free to rotate. 5. Identify a peptide bond as being cis or trans; describe why the trans orientation is favored.

5

Lecture 22 1. Use dihedral angles to describe the rotation about a single bond. 2. Describe how steric interactions cause some conformations to be higher in energy than others. 3. Explain how steric interactions between side chains in the polypeptide influence the conformation of the polypeptide chain. 4. Describe the conformation of the polypeptide chain in terms of its phi (φ), psi (ψ), and omega (ω) angles. 5. Interpret Ramachandran plots describing which phi (φ) and psi (ψ) angles are favored by amino acids in proteins. 6. Distinguish between the four levels of protein structure; be able to describe the interactions/phenomena that stabilize secondary and tertiary structure. 7. Describe the basic structural features of alpha helices and beta sheets. Lecture 23 1. Be able to identify a polynucleotide chain as being either DNA or RNA based on its chemical structure. 2. Identify the 3’ and 5’ ends of a polynucleotide chain, even if the two ends are not explicitly shown. 3. Draw the hydrogen bonds that would form between two complementary bases, given their structures. 4. Identify a DNA sequence that is complementary to a given DNA sequence. Lecture 24 1. Describe properties of gases and compare to liquids and solids. 2. Describe the key postulates of KMT and explain how KMT relates to behavior of gas particles. 3. Connect the effect of temperature on the average kinetic energies and the root-mean-square speed of gases. 4. Interpret a Boltzmann distribution of molecular speeds and match Boltzmann curves to gas samples of various temperatures and molar masses. 5. Calculate the root mean square speed of a gas at any given temperature. 6. Define effusion and diffusion, and compare effusion or diffusion rates of two gases. 7. Recognize that the pressure of the atmosphere is large. Lecture 25 1. Use the ideal gas law to calculate P, V, T, n, or mass of a gas sample given the necessary experimental information. 2. Draw and recognize graphs plotting the relationships between P, V, T, and n. 3. Know what is meant by STP. 4. Connect between density and molar mass of a gas at a given T and P.

6

Lecture 26 1. Define the partial pressure of the components of a gas mixture and determine the total pressure of the mixture. 2. Define the mole fraction of the components in a gas mixture, and relate mole fraction of a gas to its partial pressure. 3. Use ppm to describe the mole fraction of a gas. 4. Determine which gases are more likely to deviate from ideal behavior, and under which conditions. 5. Explain the effects of intermolecular forces on the behavior of a gas

Lecture 27 1. Recognize the names and definitions of the phase transitions and describe the energy changes associated with each transition. 2. Identify each phase transition (e.g., freezing, vaporization, etc.) as exothermic or endothermic. 3. Define specific heat and molar heat capacity. Know the proper units for these quantities and be able to convert between these quantities using molar mass. 4. Sketch a heating/cooling curve of a pure substance, showing how the temperature of a substance is affected by the addition/removal of heat. 5. Describe the heating curves in terms of changes in kinetic and potential energy. 6. Calculate the total amount of heat transferred when a substance changes temperature and/or undergoes a phase transition. Lecture 28 1. Define vapor pressure and describe how temperature and intermolecular forces influence vapor pressure. 2. Define the normal boiling point and predict the effect of external pressure on the boiling point of a liquid. 3. Sketch a phase diagram for a substance, including the triple point and the critical point. 4. Use a phase diagram to determine the phases of matter that will exist for a substance at specific temperatures/pressures. 5. Use a phase diagram to determine the changes that will occur to a pure substance as temperature or pressure (or both) are changed. 6. Explain how the phase diagram for water differs from most substances. 7. Relate the heating curve to the phase diagram of a pure substance. 8. Define viscosity and surface tension. 9. Given structures (formula) of molecules, predict their relative properties based on intermolecular forces (including viscosity, vapor pressure, surface tension, and enthalpy of vaporization).

7

Lecture 29 1. 2. 3. 4.

Identify types of IM forces formed and broken when a solute dissolves in a solvent. Specify factors that can make dissolution occur spontaneously. Explain, in terms of entropy, why the hydrophobic effect occurs. Explain why protein folding is favorable in terms of the enthalpic and entropic changes that take place as a protein folds. 5. Identify strong, weak, and non-electrolytes.

Lecture 30 1. Define mass fraction, weight percent, mole fraction, ppm, ppb, ppt, molarity, and convert from any concentration unit to another. 2. Calculate the concentration of ions in an electrolyte solution. 3. Determine the concentration of species (compounds or ions) after transfer or dilution. 4. Determine the concentration of species in solution when two (or more) solutions are mixed. Lecture 31 1. Define solubility and determine whether a solution is unsaturated, saturated, or supersaturated. 2. Predict whether a solute will be soluble in a solvent based on similar IMFs. 3. Describe what it means for two liquids to be miscible. Lecture 32 1. Predict the effect of temperature on solubility of solids and gases and read a plot of solubility versus temperature. 2. Predict the effect of pressure on solubility of solids and gases. 3. Use Henry’s Law to calculate the solubility of gases. Lecture 33 1. Balance simple chemical reactions. 2. Classify reactions as complete combustion, combination, decomposition, exchange, or single displacement reactions. 3. Identify the driving force for aqueous exchange reactions. 4. Predict the products for simple reactions as classified above. 5. Use the solubility guidelines to predict whether a precipitate will occur. 6. Identify spectator ions in an exchange reaction. 7. Write complete ionic reactions and net ionic reactions.

8

Lecture 34 1. Determine oxidation states of species in a reaction. 2. Determine oxidation states of atoms in organic molecules. 3. Define redox, oxidation, reduction, oxidizing agent, reducing agent, and identify these in a chemical reaction. Lecture 35 1. Use stoichiometry to determine amount of products formed or amount of reactants used. 2. Convert between moles, mass, concentration, volume, pressure, and number of atoms or molecules. 3. Determine the limiting reagent in a reaction. 4. Use limiting reagent data to determine amount of products formed or amount of excess reagent left. 5. Calculate theoretical yield and percent yield. Lecture 36 1. 2. 3. 4.

Identify the differences between RNA and DNA in terms of sugars, strandedness, and function. Describe the roles that RNA plays in the Central Dogma. Describe the three levels of hierarchical folding of RNA (primary, secondary, and tertiary). Recognize non-Watson-Crick base pairs and understand how they hydrogen bond with each other. 5. Describe the three driving forces for RNA folding. Lecture 37 1. 2. 3. 4.

Draw the products of a reaction given its arrow pushing mechanism. Explain how RNA has both genetic and functional capabilities. Describe the RNA world theory and explain the evidence that supports it. Describe how biomolecules might have arrived on early Earth or been synthesized on earth.

Lecture 38 1. 2. 3. 4.

Differentiate between a balanced chemical equation and a thermochemical equation. Draw a qualitative energy diagram for a reaction given the enthalpy change. Given the enthalpy change, identify the reaction as endothermic or exothermic. Determine the enthalpy change for a reaction given a thermochemical equation, the amount of reactants used or the amount of product formed.

Lecture 39 1. Use calorimetry data to determine the heat of reaction.

9

Lecture 40 1. Use Hess’s Law to determine the reaction enthalpy of a multistep reaction. 2. Define standard state conditions, standard states of the elements, and the standard enthalpy of formation of any element in its standard state. 3. Define the standard heat of formation and use the standard heats of formation to calculate the standard enthalpy change of any reaction; sum of products minus reactants. Lecture 41 1. Use bond dissociation energy to estimate the enthalpy change in a reaction; sum of bonds formed minus bonds broken. Lecture 42 2. Describe the basic features of antibiotic drugs, including their evolutionary origins, the major processes they inhibit, and their significance in medicine. 3. Describe the structural features of the bacterial cell wall and how it is synthesized by the cellular enzymes. 4. Explain the importance of the bacterial cell wall as it relates to osmotic pressure. Lecture 43 1. Describe the structural characteristics of β-lactam antibiotics. 2. Describe how β-lactam antibiotics react with and inhibit the function of cell wall-synthesizing enzymes. 3. Describe how β-lactamase enzymes confer resistance to β-lactams. Lecture 44 1. Describe how vancomycin inhibits cell wall synthesis via a mechanism distinct from β-lactams. 2. Describe how bacteria alter the structure of the cell wall to resist vancomycin.

10...