Identify which of the following isotopic symbols is incorrect. (6^12) C (7^13) C (12^) C
> 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
> Define each of the following radiation units: a. rad b. rem c. gray
> 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
> 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
> Describe one application of electrolytic cells.
> Will the number of isomers increase or decrease with the number of carbon atoms in a hydrocarbon? Explain your reasoning.
> Compare and contrast a battery and electrolysis.
> Write the oxidation and reduction half-reactions for the equation in Question 4.92. Question 4.92: Zn(s) + Cu2+(aq) −−−−→ Zn2+(aq) + Cu(s)
> In the following reaction, identify the oxidized species, reduced species, oxidizing agent, and reducing agent: Zn(s) + Cu2+(aq) −−−−→ Zn2+(aq) + Cu(s)
> Do metals tend to be good oxidizing agents or good reducing agents?
> During an oxidation-reduction reaction, is the oxidizing agent oxidized or reduced?
> Identify the acid and base in the following reaction: HBr (aq) + NaOH (aq) −−−−→ NaBr (aq) + H2O (l)
> During an acid-base reaction, what term is used to describe the reactant that gains a hydrogen cation, H+?
> Will a precipitate form if solutions of the soluble salts AgNO3 abd NaOH are mixed?
> Write a balanced equation for each of the following reactions: a. Nitric acid reacts with calcium hydroxide to produce water and calcium nitrate. b. Butane (C4H10) reacts with oxygen to produce water and carbon dioxide. c. Sulfur, present as an impurity
> Balance each of the following equations: a. Fe2O3(s) + CO(g) −−−−→ Fe3O4(s) + CO2(g) b. C6H6(l) + O2(g) −−−−→ CO2(g) + H2O(g) c. I4O9(s) +I2O6(s) −−−−→ I2(s) + O2(g) d. KClO3(s) −−−−→ KCl(s) + O2(g)
> Convert 300.0 K to: a. 0C b. 0F
> Complete, then balance, each of the following equations: a. Li(s) + O2(g) −−−−→ b. Ca(s) + N2(g) −−−−→ c. Al(s) + S(s) −−−−→
> Balance each of the following equations: a. C6H12O6(s) + O2(g) −−−−→ CO2(g) + H2O(g) b. H2O(l) + P4O10(s) −−−−→ H3PO4(aq) c. PCl5(g) + H2O(l) −−−−→ HCl(aq) + H3PO4(aq) d. C6H12O6(s) −−−−→ C2H6O(l) + CO2(g)
> Describe the process of checking to ensure that an equation is properly balanced.
> What is the meaning of the coefficient in a chemical equation?
> Classify each of the following reactions as decomposition (D), combination (C), single-replacement (SR), or double replacement (DR): a. KOH(s) + CO28) KHCO3(s) Δ KO(g) + CO(3) b. K,CO3(aq)- c. H2SO4(aq) + 2 NaOH(aq) d. 2AgNO3(aq) + Zn(s) 2Ag(s) + Zn
> Classify each of the following reactions as decomposition (D), combination (C), single-replacement (SR), or double-replacement (DR): a. 2Al(OH), (s)– AAl,O3(s) + 3H,O(g) b. Fe,S, (s)-A+2Fe(s) + 3S(s) c. Na,CO3(aq) + BaCl,(aq) BaCO:(s) + 2NaCl(aq) d.
> What is the meaning of (s), (l), (g), and (aq) immediately following the symbol for a chemical substance?
> What is a product? On which side of the reaction arrow are products found?
> How many mol are in 50.0 g of each of the following substances? a. Br2 b. NH4Cl c. Sr (OH)2 d. LiNO3
> How many g are required to have 0.100 mol of each of the following? a. C6H12O6 (glucose) b. NaCl c. C2H5OH (ethanol) d. Ca3(PO4)2
> For the reaction 2H2O2 (aq)−−→2H2O(l) + O2 (g) the rate law is: rate = k [H2O2] at 250C, k = 3.1 × 10-3 s-1. What effect would doubling the [H2O2] have on the rate?
> Calculate formula mass and the molar mass of CaCl2 · 2H2O.
> Calculate formula mass and the molar mass of ozone, O3.
> The formula of ascorbic acid, commonly known as vitamin C, is C6H8O6. Calculate the formula mass and molar mass of vitamin C.
> Calculate the number of carbon atoms in 15.0 g of carbon.
> What is the mass, in g, of 15.0 mol of carbon?
> Calculate the number of mol corresponding to: a. 0.10 g Ca b. 4.00 g Fe c. 2.00 kg N2
> What is the mass, in g, of 1.00 mol of nitrogen atoms?
> How many g of carbon are contained in 3.00 mol of carbon atoms?
> How many mol of sodium correspond to 1.0 × 1015 atoms of sodium?
> How many mercury atoms are present in 1.0 × 10-10 mol of mercury?
> A 4.00-g sample of Fe3O4 reacts with O2 to produce Fe2O3: 4Fe3O4 (s) 1 O2 (g)−−−−→6Fe2O3 (s) Determine the number of g of Fe2O3 produced.
> What is the mass, in g, of Avogadro’s number of iron atoms?
> What is the average molar mass of? a. S b. Na c. Hg
> How many mol of lead (Pb) atoms are equivalent to six billion lead atoms?
> What is the average mass (in amu) of? a. Zr b. Cs c. Ca
