Question: Define each of the following radiation units:

> a. Draw and name all of the isomeric products obtained from the monobromination of propane with Br2/ light. If halogenation were a completely random reaction and had an equal probability of occurring at any of the C—H bonds in a molecule, what percentage

> Predict which compound in each of the following pairs should have the higher melting and boiling points (Hint: Write the Lewis structure and determine whether the compound is ionic, polar covalent, or nonpolar covalent.): a. C2H6 and CH4 c. F2 and Br2 b

> Why do the hydrocarbon tails of membrane phospholipids provide a barrier between the inside and outside of the cell?

> How does the structure of amylose differ from that of amylopectin and glycogen?

> State Hund’s rule. Determine whether the following orbital diagrams violate Hund’s rule. а. 1s 2s 2p b. 1s 2s 2p C. 1s 2s 2p

> List the functions of several a-globulins.

> By comparing two solutions: 0.50 M sodium chloride (an ionic compound) and 0.50 M sucrose (a covalent compound). Each solution is separated from water by a semipermeable membrane. Which solution has the higher osmotic pressure?

> Distinguish among the three types of glycosphingolipids, cerebrosides, sulfatides, and gangliosides.

> Describe the purpose of a titration.

> Indicate whether the side chains of each of the amino acids listed in Question 18.1 is polar, nonpolar, basic, or acidic. Question 18.1: Write the one-letter and three-letter abbreviations, and draw the structure of each of the following amino acids. a

> What is the function of each of the four types of plasma lipoproteins?

> What is a heteropolysaccharide?

> How do antibodies protect us against infection?

> Give the Lewis structure for each of the following compounds: a. HNO3 b. CCl4 c. PBr3 d. CH3CH2OH

> How are the structural differences between triglycerides and phospholipids reflected in their different biological functions?

> An instrument for cancer treatment containing a cobalt-60 source was manufactured in 1988. In 1995, it was removed from service and, in error, was buried in a landfill with the source still in place. What percentage of its initial radioactivity will rema

> Do you think that a diet higher in omega-3 fatty acids would be an effective treatment for the symptoms of arthritis? Defend your answer.

> How would you produce a Fischer Projection beginning with a three-dimensional model of a sugar?

> Explain the bond angles of ethyne in terms of the valenceshell electron-pair repulsion (VSEPR) theory.

> Explain the bond angles of ethene in terms of the valenceshell electron-pair repulsion (VSEPR) theory.

> List the biological functions of lipids.

> Using Lewis structures and VSEPR, predict whether each of the following molecules is polar: a. CO2 b. SCl2 c. BrCl d. BCl3

> Express each of the following in units of atm: a. 10.0 torr b. 61.0 cm Hg c. 275 mm Hg d. 124 psi

> The reaction: 2NO(g) + O2 (g)−−−−→2NO2 (g) is one step in the process of forming atmospheric smog. a. How many g of NO2 can be produced by the reaction of excess O2 with 50.0 g of NO? b. If the actual yield of NO2 is 50.0 g, what is the % yield?

> Explain the mechanism by which glutamate and NO may function to promote development of memories and learning.

> For the equilibrium situation involving acetic acid, CH3COOH (aq) + H2O(l) −↽−−−−−−⇀− CH3COO-(aq) + H3O+(aq), explain the equilibrium shift occurring for the following changes: a. A strong base is added to the solution. b. More acetic acid is added to th

> Explain how the molar concentration of carbonic acid in the blood would be changed during a situation of alkalosis.

> For the reaction 47 kcal + 2SO3 (g)↽−−−−⇀2SO2 (g) + O2 (g) predict the effect on the position of the equilibrium (will it shift to the left or to the right, or will there be no change?) for each of the following changes. a. The temperature is increased.

> Where do a-amylase and b-amylase carry out their enzymatic functions?

> Titration of 17.85 mL of HNO3 solution requires 16.00 mL of 0.1600 M KOH solution. What is the molarity of the HNO3 solution?

> What metabolic defect causes galactosemia?

> How can Le Chatelier’s principle help us to increase yields of chemical reactions?

> Which of the following alcohols can be oxidized to a carboxylic acid? Name the carboxylic acid produced. For those alcohols that cannot be oxidized to a carboxylic acid, name the final product. a. Ethanol b. 2-Propanol c. 1-Propanol d. 3-Pentanol

> Describe the rules for determining the common names of carboxylic acids.

> Define the terms primary, secondary, and tertiary alcohol, and draw a general structure for each.

> For each pair of compounds, predict which would have the higher boiling point. Explain your reasoning. a. Propanamine or propanol b. Propane or ethanamine c. Methanamine or water d. Propyl methylamine or pentane

> Describe the role of disulfide bonds in the synthesis and structure of insulin.

> Calculate the osmotic pressure of 0.50 M C6H12O6 (nonelectrolyte).

> Describe the relationship between acyl group transfer and the process of protein synthesis.

> Design the synthesis of each of the following esters from organic alcohols. a. Methyl butanoate (apples) b. Octyl ethanoate (oranges)

> In your own words, describe the steps used to name a compound, using IUPAC nomenclature.

> Calculate the pH of a 1.0 × 10-5 M solution of NaOH.

> The element lithium has two naturally occurring isotopes. One of these has a mass of 6.0151 amu and a natural abundance of 7.49%. A second isotope has a mass of 7.0160 amu and a natural abundance of 92.51%. Calculate the atomic mass of lithium.

> Describe an experiment that would enable you to determine the mass (g) of solids suspended in a 1-L sample of seawater.

> Explain the significance of thioester formation in the metabolic pathways involved in fatty acid and carbohydrate breakdown.

> Cu2+(aq) can be reduced in an electrolytic cell to prepare high purity copper metal. a. Does this reduction take place at the positive or negative electrode of the electrolytic cell? Why? b. In this electrode termed the cathode or the anode? c. Write the

> Write the net ionic equation for the reaction of Pb(NO3)2(aq) with K2S(aq).

> Describe an application of reasoning involving the scientific method that has occurred in your day-to-day life.

> The tsetse fly Glossina morsitans is a large biting fly found in regions of Africa. They carry sleeping sickness, a deadly disease caused by a parasitic protozoan. The pheromone secreted by the tsetse fly contains four straight-chain alkanes: 2-methylhep

> Describe the IUPAC rules for naming cycloalkanes.

> Calculate both [H3O+] and [OH+] for a solution for which: a. pH = 5.50 b. pH = 7.00

> Consider two beakers, one containing 0.10 M NaOH and the other, 0.10 M NH3. Which solution has the greater pH? Why?

> Calculate the mass in g corresponding to: a. 0.400 mol NH3 b. 0.800 mol BaCO3 c. 2.00 mol CH4 d. 0.400 mol Ca (NO3)2

> Briefly describe the rules of the IUPAC Nomenclature System for naming ketones.

> What happens when additional solute is added to a saturated solution that is being heated?

> How does the action of glycogen phosphorylase and phosphoglucomutase result in an energy savings for the cell if the product, glucose-6- phosphate, is used directly in glycolysis?

> Write the electron configuration and shorthand electron configuration for each of the following ions: a. Rb+ b. Sr2+ c. S2- d. I-

> What volume of 0.2000 M KOH is required to titrate 25.00 mL of 0.1500 M HNO3?

> What points of Dalton’s theory are no longer current?

> Determine the number of protons and electrons in each of the following ions. a. S2- b. K+ c. Cd2+

> a. Explain why a negative ion is always larger than its parent atom. b. Explain why a sodium ion is commonly found in nature but a sodium atom is not.

> The following diagram represents the endothermic reaction at equilibrium at 258C: heat + A(g) ↽−−−⇀2B(g). a. The temperature is increased to 500C. Which of the

> In Question 1.123, you calculated the volume of 8.00 × 102 g of air with a density of 1.29 g/L. The temperature of the air sample was lowered and the density increased to 1.50 g/L. Calculate the new volume of the air sample. Question 1.123: What volume

> a. Describe the process of fusion. b. How could this process be used for the production of electrical energy?

> How is the negative charge of a polyatomic anion incorporated when determining the number of valence electrons to be used in a Lewis structure?

> When drawing a Lewis structure, which elements can never be a central atom?

> Would H2O or CCl4 be expected to have a higher melting point? Why?

> How many total electrons and valence electrons are found in an atom of each of the following elements? What is the number of the principal energy level in which the valence electrons are found? a. Mg b. K c. C d. Br e. Ar f. Xe

> Draw the Lewis structures and shapes of each of the following molecules. Identify the molecular geometry of each shape using VSEPR. a. C2H4 b. C2H2

> Which form of radiation has the longer wavelength, ultraviolet or infrared? Explain your reasoning.

> Contrast ionic and covalent compounds with respect to their behaviors in solution.

> Write a suitable formula for: a. diphosphorus pentoxide b. dioxygen difluoride

> Write a suitable formula for: a. ammonium acetate b. ammonium cyanide

> Write a suitable formula for: a. aluminum nitrate b. potassium nitrate

> Write the correct formula for each of the following: a. magnesium carbonate b. magnesium bicarbonate

> Write the correct formula for each of the following: a. potassium oxide b. potassium nitride

> Name each of the following compounds: a. Na2O b. Fe (OH)3 c. CaBr2

> How do orbits and orbitals differ?

> If the % yield of Fe2O3 in Question 4.116 is 90.0%, what is the actual yield of Fe2O3? Question 4.116: A 4.00-g sample of Fe3O4 reacts with O2 to produce Fe2O3: 4Fe3O4 (s) 1 O2 (g)−−−−→6Fe2O3 (s)

> Determine the number of protons and electrons in each of the following ions: a. Ni2+ b. Br- c. N3-

> Predict the formula of a compound formed from: a. boron and hydrogen b. magnesium and phosphorus

> Write the formula for each of the following polyatomic ions: a. the phosphate ion b. the cyanide ion

> Write the formula for each of the following monatomic ions: a. the calcium ion b. the chromium (VI) ion

> Name each of the following ions: a. ClO- b. NH4+ c. CH3COO-

> Predict the formula of the compound formed from the combination of ions of calcium and fluorine.

> Explain, using Lewis symbols and the octet rule, why neon is so nonreactive.

> When there is a reaction between each of these pairs of atoms, ions form. Using Lewis symbols, write the reactions showing how electrons are lost or gained when these atoms become ions. a. Na + O b. Na + S c. Si + H

> Identify which of the following isotopic symbols is incorrect. (6^12) C (7^13) C (12^) C

> Describe the difference between nonpolar covalent and polar covalent bonding.

> Draw the appropriate Lewis symbol for each of the following ions: a. Be2+ b. Al3+ c. O2- d. S2-

> Arrange each of the following lists of elements in order of increasing atomic size: a. Al, Si, P b. In, Ga, Al c. Sr, Ca, Ba d. P, N, Sb

> Round each of the following numbers to two significant figures. a. 6.2262 b. 3895 c. 6.885 d. 2.2247 e. 0.0004109

> In each of the following pairs of compounds, choose the compound with the higher melting point. a. KF and F2 b. CO and O2 c. NBr3 and CCl4

> Would you expect ethylamine (Question 3.84) to dissolve in water? Question 3.84: Ethylamine is an example of an important class of organic compounds. The molecular formula of ethylamine is CH3CH2NH2.

> What effect does polarity have on the boiling point of a pure compound?

> Using the VSEPR theory, predict the geometry, polarity, and water solubility of each compound in Question 3.80. Question 3.80: a. HNO3 b. CCl4 c. PBr3 d. CH3CH2OH

> What effect does polarity have on the solubility of a compound in water?

> Which of the following compounds have polar bonds but are nonpolar covalent compounds? a. SO2 b. CF4 c. NH3

> Do nonmetals tend to gain or lose electrons? Do they become cations or anions?

> For each of the following element symbols, give the name of the element, its atomic number, and its atomic mass. a. Mg b. Ne c. Se