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?
> 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
> 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
> 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
> Identify the oxidizing agent, reducing agent, substance oxidized, and substance reduced in the reaction described in Question 4.10. Question 4.10: Write the oxidation half-reaction, the reduction half-reaction, and the complete reaction for the formati
> If the actual yield of oxygen gas in Question 4.114 is 1.10 × 10-2 g, what is the % yield? Question 4.114: Chemical Control of Microbes (Section 4.8) describes the breakdown of the antiseptic H2O2 with the balanced equation 2H2O2(aq) −−−−→ 2H2O(l) + O2
> The reaction of calcium hydride with water can be used to prepare hydrogen gas: CaH2 (s) 1 2H2O(l)−−−→Ca (OH)2 (aq) + 2H2 (g) How many g of hydrogen gas are produced in the reaction of 1.00 × 102 g calcium hydride with water?
> Dinitrogen monoxide (also known as nitrous oxide and used as an anesthetic) can be made by heating ammonium nitrate: NH4NO3 (s) −−D−→ N2O(g) + 2H2O(g) How many g of dinitrogen monoxide can be made from 1.00 3 102 g of ammonium nitrate?
> A 3.5-g sample of water reacts with PCl3 according to the following equation: 3H2O(l) + PCl3 (g)−−−→H3PO3 (aq) + 3HCl(aq) How many mol of H3PO3 are produced?
> How many g of Al will react with 3.00 mol of O2? 4Al(s) + 3O2 (g)−−−→2Al2O3 (s)
> Chemical Control of Microbes (Section 4.8) describes the breakdown of the antiseptic H2O2 with the balanced equation 2H2O2(aq) −−−−→ 2H2O(l) + O2(g). Assuming there is an unlimited amount of the enzyme, how many g of O2 would be produced from 1.00 × 10-1
> Describe the steps used in the calculation of g of product resulting from the reaction of a specified number of g of reactant.
> Write the oxidation half-reaction, the reduction half-reaction, and the complete reaction for the formation of calcium iodide from calcium metal and I2. Remember, the electron gain must equal the electron loss.
> Write the two conversion factors that can be written for the relationship between cm and in.
> The following four measurements were made for an object whose true volume is 17.55 mL. 18.69 mL, 18.69 mL, 18.70 mL, 18.71 mL Describe the measurements in terms of their accuracy and their precision.
> For any given principal energy level, what is the maximum number of electrons that can exist in the following subshells? a. s b. p c. d
> Distinguish between a sublevel and an orbital.
> Use the graph prepared in Question 2.75 to predict the melting point of francium (Fr). Question 2.75: Element Atomic Number Melting Point (C) Li 3 180.5 Na 11 97.8 K 19 63.3 Rb 37 38.9 Cs 55 28.4
> Which group of the periodic table is known as the alkaline earth metals? List their symbols.
> What English unit of volume is similar to a L?
> What was the major deficiency of Bohr’s atomic model?
> What is the relationship between density and specific gravity?
> Give two reasons why the Bohr theory did not stand the test of time.
> Assume the Body-Mass Index (BMI) is calculated using the expression BMI = weight (kg)/height2 (m2). If a patient has a height of 1.6 m and a BMI of 38 kg/m2, what is the patient’s weight in both kg and lb?
> Critique this statement: Promotion of electrons is accompanied by a release of energy.
> The density of mercury is 13.6 g/mL. If a sample of mercury weighs 272 g, what is the volume of the sample in mL?
> When electrical energy is applied to an element in its gaseous state, light is produced. How does the light differ among elements?
> The specific gravity of a patient’s urine sample was measured to be 1.008. Given that the density of water is 1.000 g/mL at 40C, what is the density of the urine sample?
> What is the relationship between the energy of light and its wavelength?
> The density of methanol at 200C is 0.791 g/mL. What is the mass of a 50.0 mL sample of methanol?
> Describe electromagnetic radiation according to its particle nature.
> Refer to Question 1.129. Suppose that each of the bars had the same mass. How could you determine which bar had the lowest density and which had the highest density? Question 1.129: You are given three bars of metal. Each is labeled with its identity (
> Predict the polarity of each compound in Question 3.98. Question 3.98: a. SeO2 b. SeO3
> You are given a piece of wood that is either maple, teak, or oak. The piece of wood has a volume of 1.00 × 102 cm3 and a mass of 98 g. The densities of maple, teak, and oak are as follows: What is the identity of the piece of wood? Woo
> K+ and Ar are isoelectronic. Which is larger? Why?
> What is meant by the term electromagnetic spectrum?
> What is the mass of a femur (leg bone) having a volume of 118 cm3? The density of bone is 1.8 g/cm3.
> Draw a diagram representing a homogeneous mixture of two different substances. Use two different colored spheres to represent the two different substances.
> Arrange each of the following lists of elements in order of decreasing electron affinity: a. Mg, P, Cl b. Br, I, Cl
> Calculate the density of 50.0 g of an isopropyl alcohol– water mixture (commercial rubbing alcohol) that has a volume of 63.6 mL.
