Question: In general, how do we envision the

> What is an element? Which elements are most abundant on the earth? Which elements are most abundant in the human body?

> What are proteins? Are proteins polymers? Explain. What molar mass ranges are shown by proteins? What fraction of our bodies is made up of proteins?

> What are trace elements, and why are such elements important to the body’s health?

> is the study of the chemistry of living systems.

> Complete and balance the following equations. a. Pb(NO3)2(aq) + Na2S(aq) / b. AgNO3(aq) + HCl(aq) / c. 2Mg(s) + O2(g) / d. H2SO4(aq) + 2KOH(aq) / e. BaCl2(aq) + H2SO4(aq) / f. Mg(s) + H2SO4(aq) / g. 2Na3PO3(aq) + 3CaCl2(aq) / h. 2C4H10(l) + 13O2(g) /

> Write the balanced net ionic equation for the reaction that takes place when aqueous solutions of the following solutes are mixed. If no reaction is likely, explain why no reaction would be expected for that combination of solutes. a. potassium nitrate

> What is meant by denaturation of a protein? Give three examples of situations in which proteins are denatured.

> Give balanced equations for two examples of each of the following types of reactions. a. precipitation b. single-displacement c. combustion d. synthesis e. oxidation – reduction f. decomposition g. acid–base neutralization

> Balance each equation. Which equations can be classified as oxidation–reduction reactions? a. FeO(s) + HNO3(aq) Fe(NO3)2(aq) + H2O(l) b. Mg(s) + CO2(g) + O2(g) MgCO3(s) c. NaOH(s) + CuSO4(aq) Cu(OH)2(s) + N

> The reagent shelf in a general chemistry lab contains aqueous solutions of the following substances: silver nitrate, sodium chloride, acetic acid, nitric acid, sulfuric acid, potassium chromate, barium nitrate, phosphoric acid, hydrochloric acid, lead ni

> Balance each of the following chemical equations. a. Na2SO4(aq) + BaCl2(aq) BaSO4(s) + NaCl(aq) b. Zn(s) + H2O(g) ZnO(s) + H2(g) c. NaOH(aq) + H3PO4(aq) Na3PO4(aq) + H2O(l) d. Al(s) + Mn2O3(s)

> List and define all the ways of classifying chemical reactions that have been discussed in the text. Give a balanced chemical equation as an example of each type of reaction, and show clearly how your example fits the definition you have given.

> Give an example of a synthesis reaction and of a decomposition reaction. Are synthesis and decomposition reactions always also oxidation–reduction reactions? Explain.

> What is a combustion reaction? Are combustion reactions a unique type of reaction, or are they a special case of a more general type of reaction? Write an equation that illustrates a combustion reaction.

> What do we call reactions in which electrons are transferred between atoms or ions? What do we call a loss of electrons by an atom or ion? What is it called when an atom or ion gains electrons? Can we have a process in which electrons are lost by one spe

> What is a salt? How are salts formed by acid–base reactions? Write chemical equations showing the formation of three different salts. What other product is formed when an aqueous acid reacts with an aqueous base? Write the net ionic equation for the form

> Describe some physical and chemical properties of acids and bases. What is meant by a strong acid or base? Are strong acids and bases also strong electrolytes? Give several examples of strong acids and strong bases.

> a. Given that 100. mL of ethyl alcohol weighs 78.5 g, calculate the density of ethyl alcohol. b. What volume would 1.59 kg of ethyl alcohol occupy? c. What is the mass of 1.35 L of ethyl alcohol? d. Pure aluminum metal has a density of 2.70 g/cm3. Cal

> In general terms, what are the spectator ions in a precipitation reaction? Why are the spectator ions not included in writing the net ionic equation for a precipitation reaction? Does this mean that the spectator ions do not have to be present in the sol

> Summarize the simple solubility rules for ionic compounds. How do we use these rules in determining the identity of the solid formed in a precipitation reaction? Give examples including balanced complete and net ionic equations.

> Define the term strong electrolyte. What types of substances tend to be strong electrolytes? What does a solution of a strong electrolyte contain? Give a way to determine if a substance is a strong electrolyte.

> Explain to your friend what chemists mean by a precipitation reaction. What is the driving force in a precipitation reaction? Using the information provided about solubility in these chapters, write balanced molecular and net ionic equations for five exa

> What is meant by the driving force for a reaction? Give some examples of driving forces that make reactants tend to form products. Write a balanced chemical equation illustrating each type of driving force you have named.

> When balancing a chemical equation, why is it not permissible to adjust the subscripts in the formulas of the reactants and products? What would changing the subscripts within a formula do? What do the coefficients in a balanced chemical equation represe

> What does it mean to “balance” an equation? Why is it so important that equations be balanced? What does it mean to say that atoms must be conserved in a balanced chemical equation? How are the physical states of reactants and products indicated when wri

> What, in general terms, does a chemical equation indicate? What are the substances indicated to the left of the arrow called in a chemical equation? To the right of the arrow?

> What kind of visual evidence indicates that a chemical reaction has occurred? Give an example of each type of evidence you have mentioned. Do all reactions produce visual evidence that they have taken place?

> a. Fill in the following table as if it is a well plate and you are mixing two aqueous compounds at a time to see if a precipitate forms. If a precipitate is expected to form, indicate that by writing the correct formula for the precipitate in the corres

> In general terms, what does the tertiary structure of a protein describe? Clearly distinguish between the secondary and tertiary structures.

> Which of the following solutions contains the greatest number of ions? a. 100.0 mL of 1.0 M sodium nitrate b. 100.0 mL of 1.0 M iron(III) nitrate c. 100.0 mL of 1.0 M copper(II) nitrate d. 100.0 mL of 1.0 M calcium nitrate e. All of the solutions ab

> In an open flask, 20.0 mL of an aqueous solution (density of solution = 1.103 g/mL) is combined with 13.5 g of a solid, and a chemical reaction takes place. One of the reaction products is 1.473 L of gas with density 5 1.798 g/L. What is the mass of the

> Rank the following species from lowest to highest boiling point: N2(l), Ne(l), BeO(l), CO(l)

> If 125 mL of concentrated sulfuric acid solution (density 1.84 g/mL, 98.3% H2SO4 by mass) is diluted to a final volume of 3.01 L, calculate the following information. a. the mass of pure H2SO4 in the 125-mL sample. b. the molarity of the concentrated a

> Calculate the volume (in milliliters) of each of the following acid solutions that would be required to neutralize 36.2 mL of 0.259 M NaOH solution. a. 0.271 M HCl b. 0.119 M H2SO4 c. 0.171 M H3PO4

> Chlorine gas, Cl2, can be generated in small quantities by addition of concentrated hydrochloric acid to manganese(IV) oxide. MnO2(s) + 4HCl(aq) / MnCl2(aq) + 2H2O(l) + Cl2(g) The chlorine gas is bubbled through water to dissolve any traces of HCl remai

> Calculate the indicated quantity for each gas sample. a. The volume occupied by 1.15 g of helium gas at 25 °C and 1.01 atm pressure. b. The partial pressure of each gas if 2.27 g of H2 and 1.03 g of He are confined to a 5.00-L container at 0 °C. c. Th

> What is one equivalent of an acid? What does an equivalent of a base represent? How is the equivalent weight of an acid or a base related to the substance’s molar mass? Give an example of an acid and a base that have equivalent weights equal to their mol

> When a solution is diluted by adding additional solvent, the concentration of solute changes but the amount of solute present does not change. Explain. Suppose 250. mL of water is added to 125 mL of 0.551 M NaCl solution. Explain how you would calculate

> Define a saturated solution. Does saturated mean the same thing as saying the solution is concentrated? Explain. Why does a solute dissolve only to a particular extent in water? How does formation of a saturated solution represent an equilibrium?

> Without performing the actual calculations, determine to how many significant figures the results of the following calculations should be reported. c. 1.782 + 0.00035 + 2.11 d. (6.521)(5.338 + 2.11) e. 9 - 0.000017 f. (4.2005 * 2.7)(7.99118) g. (5.12

> Define a solution. Describe how an ionic solute such as NaCl dissolves in water to form a solution. How are the strong bonding forces in a crystal of ionic solute overcome? Why do the ions in a solution not attract each other so strongly as to reconstitu

> Define the bonding that exists in metals and how this model explains some of the unique physical properties of metals. What are metal alloys? Identify the two main types of alloys, and describe how their structures differ. Give several examples of each t

> Define a crystalline solid. Describe in detail some important types of crystalline solids and name a substance that is an example of each type of solid. Explain how the particles are held together in each type of solid (the interparticle forces that exis

> Why does the process of vaporization require an input of energy? Why is it so important that water has a large heat of vaporization? What is condensation? Explain how the processes of vaporization and condensation represent an equilibrium in a closed con

> Define London dispersion forces. Draw a picture showing how London forces arise. Are London forces relatively strong or relatively weak? Explain. Although London forces exist among all molecules, for what type of molecule are they the only major intermol

> What is a dipole–dipole attraction? How do the strengths of dipole–dipole forces compare with the strengths of typical covalent bonds? What is hydrogen bonding? What conditions are necessary for hydrogen bonding to exist in a substance or mixture? What e

> Are changes in state physical or chemical changes? Explain. What type of forces must be overcome to melt or vaporize a substance (are these forces intramolecular or intermolecular)? Define the molar heat of fusion and molar heat of vaporization. Why is t

> Define the normal boiling point of water. Why does a sample of boiling water remain at the same temperature until all the water has been boiled? Define the normal freezing point of water. Sketch a representation of a heating/cooling curve for water, mark

> Describe some of the physical properties of water. Why is water one of the most important substances on earth?

> For each of the following, make the indicated conversion, showing explicitly the conversion factor(s) you used. a. 593.2 kg to grams b. 593.2 lbs to grams c. 8.312 km to miles d. 8.312 ft to miles e. 6.219 ft to meters f. 6.219 cm to meters g. 329

> What does “STP” stand for? What conditions correspond to STP? What is the volume occupied by one mole of an ideal gas at STP?

> Without consulting your textbook, list and explain the main postulates of the kinetic molecular theory for gases. How do these postulates help us account for the following bulk properties of a gas: the pressure of the gas and why the pressure of the gas

> Dalton’s law of partial pressures concerns the properties of mixtures of gases. What is meant by the partial pressure of an individual gas in a mixture? How does the total pressure of a gaseous mixture depend on the partial pressures of the individual ga

> What do we mean specifically by an ideal gas? Explain why the ideal gas law (PV = nRT) is actually a combination of Boyle’s, Charles’s, and Avogadro’s gas laws. What is the numerical value and what are the specific units of the universal gas constant, R?

> What does Avogadro’s law tell us about the relationship between the volume of a sample of gas and the number of molecules the gas contains? Why must the temperature and pressure be held constant for valid comparisons using Avogadro’s law? Does Avogadro’s

> Explain how the concept of absolute zero came about through Charles’s studies of gases. Hint: What would happen to the volume of a gas sample at absolute zero (if the gas did not liquefy first)? What temperature scale is defined with its lowest point as

> What does Charles’s law tell us about how the volume of a gas sample varies as the temperature of the sample is changed? How does this volume–temperature relationship differ from the volume–pressure relationship of Boyle’s law? Give two mathematical expr

> When using Boyle’s law in solving problems in the textbook, you may have noticed that questions were often qualified by stating that “the temperature and amount of gas remain the same.” Why was this qualification necessary?

> Your textbook gives several definitions and formulas for Boyle’s law for gases. Write, in your own words, what this law really tells us about gases. Now write two mathematical expressions that describe Boyle’s law. Do these two expressions tell us differ

> What is the SI unit of pressure? What units of pressure are commonly used in the United States? Why are these common units more convenient to use than the SI unit? Describe a manometer and explain how such a device can be used to measure the pressure of

> How is the secondary structure of a protein related to its function in the body? Give examples.

> How does the pressure of the atmosphere arise? Sketch a representation of the device commonly used to measure the pressure of the atmosphere. Your textbook described a simple experiment to demonstrate the pressure of the atmosphere. Explain this experime

> What are some of the general properties of gases that distinguish them from liquids and solids?

> Describe a buffered solution. Give three examples of buffered solutions. For each of your examples, write equations and explain how the components of the buffered solution consume added strong acids or bases. Why is buffering of solutions in biological s

> The solubility product of magnesium carbonate, MgCO3, has the value Ksp = 6.82 * 10-6 at 25 0C. How many grams of MgCO3 will dissolve in 1.00 L of water?

> Calculate the pH and pOH values for each of the following solutions. a. 0.00562 M HClO4 b. 3.98 * 10-4 M KOH c. 0.078 M HNO3 d. 4.71 * 10-6 M Ca(OH)2

> For each of the following, calculate the indicated quantity. a. [OH-] = 2.11 * 10-4 M, [H+] =? b. [OH2] = 7.34 * 10-6 M, pH 5 ? c. [OH2] = 9.81 * 10-8 M, pOH 5 ? d. pH = 9.32, pOH =? e. [H+] = 5.87 * 10-11 M, pH =? f. pH = 5.83, [H+] =?

> Explain how dissolving a slightly soluble salt to form a saturated solution is an equilibrium process. Give three balanced chemical equations for solubility processes and write the expressions for Ksp corresponding to the reactions you have chosen. When

> In your own words, paraphrase Le Châtelier’s principle. Give an example (including a balanced chemical equation) of how each of the following changes can affect the position of equilibrium in favor of additional products for a system: the concentration o

> Compare homogeneous and heterogeneous equilibria. Give a balanced chemical equation and write the corresponding equilibrium constant expression as an example of each of these cases. How does the fact that an equilibrium is heterogeneous influence the exp

> Although the equilibrium constant for a given reaction always has the same value at the same temperature, the actual concentrations present at equilibrium may differ from one experiment to another. Explain. What do we mean by an equilibrium position? Is

> What is a mixture? What is a solution? How do mixtures differ from pure substances? What are some of the techniques by which mixtures can be resolved into their components?

> Describe how we write the equilibrium expression for a reaction. Give three examples of balanced chemical equations and the corresponding expressions for their equilibrium constants.

> Sketch a graph for the progress of a reaction illustrating the activation energy for the reaction. Define “activation energy.” Explain how an increase in temperature for a reaction affects the number of collisions that possess an energy greater than Ea.

> Explain the collision model for chemical reactions. What “collides”? Do all collisions result in the breaking of bonds and formation of products? Why? How does the collision model explain why higher concentrations and higher temperatures tend to make rea

> How is the pH scale defined? What range of pH values corresponds to acidic solutions? What range corresponds to basic solutions? Why is pH = 7.00 considered neutral? When the pH of a solution changes by one unit, by what factor does the hydrogen ion conc

> Explain how water is an amphoteric substance. Write the chemical equation for the autoionization of water. Write the expression for the equilibrium constant, Kw, for this reaction. What values does Kw have at 25 0C? What are [H+] and [OH-] in pure water

> How is the strength of an acid related to the position of its ionization equilibrium? Write the equations for the dissociation (ionization) of HCl, HNO3, and HClO4 in water. Since all these acids are strong acids, what does this indicate about the basici

> Acetic acid is a weak acid in water. What does this indicate about the affinity of the acetate ion for protons compared to the affinity of water molecules for protons? If a solution of sodium acetate is dissolved in water, the solution is basic. Explain.

> Describe the relationship between a conjugate acid–base pair in the Brønsted–Lowry model. Write balanced chemical equations showing the following molecules/ions behaving as Brønsted– Lowry acids in water: HCl, H2SO4, H3PO4, NH4+. Write balanced chemical

> How are the Arrhenius and Brønsted–Lowry definitions of acids and bases similar, and how do these definitions differ? Could a substance be an Arrhenius acid but not a Brønsted– Lowry acid? Could a substance be a Brønsted–Lowry acid but not an Arrhenius a

> What is an element, and what is a compound? Give examples of each. What does it mean to say that a compound has a constant composition? Would samples of a particular compound here and in another part of the world have the same composition and properties?

> Hydrogen gas and oxygen gas react violently to form water. When this occurs, a very loud noise is heard. a. Draw the Lewis structures for hydrogen gas, oxygen gas, and water. b. State whether each molecule is polar or nonpolar and why. Explain how the p

> Draw the Lewis structure for each of the following molecules or ions. Indicate the number and spatial orientation of the electron pairs around the boldface atom in each formula. Predict the simple geometric structure of each molecule or ion, and indicate

> Based on the electron configuration of the simple ions that the following pairs of elements would be expected to form, predict the formula of the simple binary compound that would be formed by each pair. a. Al and Cl b. Na and N c. Mg and S d. Ca and

> Which of the following has the smallest ionization energy? a. Se2- b. Br- c. Sr2+ d. Zr4+ e. Rb+

> An unknown element is a nonmetal and has a valence electron configuration of ns2np4. a. How many valence electrons does this element have? b. Which of the following are possible identities for this element? Cl, S, Pb, Se, Cr c. What is the general for

> What do we mean by the geometric structure of a molecule? Draw the geometric structures of at least four simple molecules of your choosing and indicate the bond angles in the structures. Explain the main ideas of the valence shell electron pair repulsion

> Although many simple molecules fulfill the octet rule, some common molecules are exceptions to this rule. Give three examples of molecules whose Lewis structures are exceptions to the octet rule.

> What does a double bond between two atoms represent in terms of the number of electrons shared? What does a triple bond represent? When writing a Lewis structure, explain how we recognize when a molecule must contain double or triple bonds. What are reso

> For three simple molecules of your own choice, apply the rules for writing Lewis structures. Write your discussion as if you are explaining the method to someone who is not familiar with Lewis structures.

> In writing Lewis structures for molecules, what is meant by the duet rule? To which element does the duet rule apply? What do we mean by the octet rule? Why is attaining an octet of electrons important for an atom when it forms bonds to other atoms? What

> It is important to be able to distinguish between the physical and the chemical properties of chemical substances. Choose a chemical substance you are familiar with, then use the Internet or a handbook of chemical information to list three physical prope

> Why does a Lewis structure for a molecule show only the valence electrons? What is the most important factor for the formation of a stable compound? How do we use this requirement when writing Lewis structures?

> Give evidence that ionic bonds are very strong. Does an ionic substance contain discrete molecules? With what general type of structure do ionic compounds occur? Sketch a representation of a general structure for an ionic compound. Why is a cation always

> How is the attainment of a noble gas electron configuration important to our ideas of how atoms bond to each other? When atoms of a metal react with atoms of a nonmetal, what type of electron configurations do the resulting ions attain? Explain how the a

> What does it mean to say that a molecule has a dipole moment? What is the difference between a polar bond and a polar molecule (one that has a dipole moment)? Give an example of a molecule that has polar bonds and that has a dipole moment. Give an exampl

> What is meant by electronegativity? How is the difference in electronegativity between two bonded atoms related to the polarity of the bond? Using Fig. 12.3, give an example of a bond that would be nonpolar and of a bond that would be highly polar.

> In the formation of a polynucleotide (a short portion of the DNA molecule), which components (sugar, base, or phosphate) on adjacent nucleotides bond to each other?

> What do we mean by covalent bonding and polar covalent bonding? How are these two bonding types similar and how do they differ? What circumstance must exist for a bond to be purely covalent? How does a polar covalent bond differ from an ionic bond?