Question: According to Bohr, what happens to the

> Use the periodic table shown in Fig. 4.9 to determine the atomic mass (per mole) or molar mass of each of the substances in column 1, and find that mass in column 2. Column 1………â€&b

> A compound containing only sulfur and nitrogen is 69.6% S by mass. The molar mass is 184 g/mol. What is the correct name for this compound?

> Identify each of the following unbalanced reaction equations as belonging to one or more of the following categories: precipitation, acid–base, or oxidation–reduction. a. Fe(s) + H2SO4(aq) Fe3(SO4)2(aq) + H2(g) b. HClO4(aq) + RbOH(

> Balance the equation for each of the following oxidation– reduction chemical reactions. a. Na(s) + O2(g) Na2O2(s) b. Fe(s) + H2SO4(aq) FeSO4(aq) + H2(g) c. Al2O3(s) Al(s) + O2(g) d. Fe(s) + Br2(l

> If you spilled a cup of freshly brewed hot tea on yourself, you would be burned. If you spilled the same quantity of iced tea on yourself, you would not be burned. Explain.

> Using the average atomic masses given inside the front cover of the text, calculate how many moles of each substance the following masses represent. a. 4.21 g of copper(II) sulfate b. 7.94 g of barium nitrate c. 1.24 mg of water d. 9.79 g of tungsten

> A compound has been analyzed and has been found to have the following composition: copper, 66.75%; phosphorus, 10.84%; oxygen, 22.41%. Determine the empirical formula of the compound.

> A binary compound of boron and hydrogen has the following percentage composition: 78.14% boron, 21.86% hydrogen. If the molar mass of the compound is determined by experiment to be between 27 and 28 g, what are the empirical and molecular formulas of the

> A compound used in the nuclear industry has the following composition: uranium, 67.61%; fluorine, 32.39%. Determine the empirical formula of the compound.

> How does the molecular formula of a compound differ from the empirical formula? Can a compound’s empirical and molecular formulas be the same? Explain.

> Phosphorus and chlorine form two binary compounds, in which the percentages of phosphorus are 22.55% and 14.87%, respectively. Calculate the empirical formulas of the two binary phosphorus–chlorine compounds.

> On the basis of the general solubility rules given in Table 7.1, predict the identity of the precipitate that forms when aqueous solutions of the following substances are mixed. If no precipitate is likely, indicate why (which rules apply). a. iron(III)

> Tetraphenylporphyrin is a synthetic compound that resembles naturally occurring porphyrins. Porphyrins are found in hemoglobin and cytochromes, which are important to biological functions in humans. Tetraphenylporphyrin is composed of only C, H, and N at

> For the cations listed in the left-hand column, give the formulas of the precipitates that would form with each of the anions in the right-hand column. If no precipitate is expected for a particular combination, so indicate. Cations ……….... Anions Ag+…

> When lithium metal is heated strongly in an atmosphere of pure nitrogen, the product contains 59.78% Li and 40.22% N on a mass basis. Determine the empirical formula of the compound.

> For each of the following reactions, give the balanced equation for the reaction and state the meaning of the equation in terms of the numbers of individual molecules and in terms of moles of molecules. a. PCl3(l) + H2O(l) H3P

> A compound has the following percentage composition by mass: copper, 33.88%; nitrogen, 14.94%; oxygen, 51.18%. Determine the empirical formula of the compound.

> A compound was analyzed and was found to contain the following percentages of the elements by mass: lithium, 46.46%; oxygen, 53.54%. Determine the empirical formula of the compound.

> Distinguish between the molecular equation, the complete ionic equation, and the net ionic equation for a reaction in solution. Which type of equation most clearly shows the species that actually react with one another?

> If 2.50 g of aluminum metal is heated in a stream of fluorine gas, it is found that 5.28 g of fluorine will combine with the aluminum. Determine the empirical formula of the compound that results.

> If 1.25 g of aluminum metal is heated in an atmosphere of fluorine gas, 3.89 g of aluminum fluoride results. Determine the empirical formula of aluminum fluoride.

> If cobalt metal is mixed with excess sulfur and heated strongly, a sulfide is produced that contains 55.06% cobalt by mass. Calculate the empirical formula of the sulfide.

> A compound was analyzed and was found to contain the following percentages of the elements by mass: tin, 45.56%; chlorine, 54.43%. Determine the empirical formula of the compound.

> By now, you are familiar with enough chemical compounds to begin to write your own chemical reaction equations. Write two examples each of what we mean by a synthesis reaction and by a decomposition reaction.

> If a 1.271-g sample of aluminum metal is heated in a chlorine gas atmosphere, the mass of aluminum chloride produced is 6.280 g. Calculate the empirical formula of aluminum chloride.

> A compound was analyzed and was found to contain the following percentages of the elements by mass: boron, 78.14%; hydrogen, 21.86%. Determine the empirical formula of the compound.

> For each of the following reactions, give the balanced chemical equation for the reaction and state the meaning of the equation in terms of individual molecules and in terms of moles of molecules. a. MnO2(s) + Al(s) Mn(s) +

> A 0.5998-g sample of a new compound has been analyzed and found to contain the following masses of elements: carbon, 0.2322 g; hydrogen, 0.05848 g; oxygen, 0.3091 g. Calculate the empirical formula of the compound.

> A compound was analyzed and was found to contain the following percentages of the elements by mass: nitrogen, 11.64%; chlorine, 88.36%. Determine the empirical formula of the compound.

> A compound was analyzed and was found to contain the following percentages of the elements by mass: barium, 89.56%; oxygen, 10.44%. Determine the empirical formula of the compound.

> Although many sulfate salts are soluble in water, calcium sulfate is not (Table 7.1). Therefore, a solution of calcium chloride will react with sodium sulfate solution to produce a precipitate of calcium sulfate. The balanced equation is CaCl2(aq) + Na2

> A common method for determining how much chloride ion is present in a sample is to precipitate the chloride from an aqueous solution of the sample with silver nitrate solution and then to weigh the silver chloride that results. The balanced net ionic rea

> If steel wool (iron) is heated until it glows and is placed in a bottle containing pure oxygen, the iron reacts spectacularly to produce iron(III) oxide. Fe(s) + O2(g) Fe2O3(s) If 1.25 g of iron is heated and placed in a bottle co

> Lead(II) oxide from an ore can be reduced to elemental lead by heating in a furnace with carbon. PbO(s) + C(s) Pb(l) + CO(g) Calculate the expected yield of lead if 50.0 kg of lead oxide is heated with 50.0 kg of carbon.

> For each of the following ionic substances, calculate the percentage of the overall molar mass of the compound that is represented by the negative ions in the substance. a. ammonium sulfide b. calcium chloride c. barium oxide d. nickel(II) sulfate

> For each of the following samples of ionic substances, calculate the number of moles and mass of the positive ions present in each sample. a. 4.25 g of ammonium iodide, NH4I b. 6.31 moles of ammonium sulfide, (NH4)2S c. 9.71 g of barium phosphide, Ba3

> What is the mass percent of oxygen in each of the following compounds? a. carbon dioxide b. sodium nitrate c. iron(III) phosphate d. ammonium carbonate e. aluminum sulfate

> Using the average atomic masses for each of the following elements (see the table inside the cover of this book), calculate the mass (in amu) of each of the following samples. a. 125 carbon atoms b. 5 million potassium atoms c. 1.04 * 1022 lithium ato

> Calculate the percent by mass of the element listed first in the formulas for each of the following compounds. a. adipic acid, C6H10O4 b. ammonium nitrate, NH4NO3 c. caffeine, C8H10N4O2 d. chlorine dioxide, ClO2 e. cyclohexanol, C6H11OH f. dextrose

> Complete the following table. Value of n Possible Sublevels 1 3 4

> When the electron in hydrogen is in the n = 3 principal energy level, the atom is in a/an state.

> The higher the principal energy level, n, the (closer to/farther from) the nucleus is the electron.

> What are the differences between the 2s orbital and the 1s orbital of hydrogen? How are they similar?

> Consider the following representation of a set of p orbitals for an atom: Which of the following statements is(are) true? a. The areas represented by the p orbitals are positively charged clouds with negatively charged electrons embedded within these

> What is electromagnetic radiation? At what speed does electromagnetic radiation travel?

> In the 1920s, de Broglie and Schrödinger developed what is now called wave mechanics or quantum mechanics. Which of the following statements with regard to this model is(are) true? a. The position of an electron can be exactly found and measured. b. Wit

> Calculate the percent by mass of the element listed first in the formulas for each of the following compounds. a. copper(II) bromide, CuBr2 b. copper(I) bromide, CuBr c. iron(II) chloride, FeCl2 d. iron(III) chloride, FeCl3 e. cobalt(II) iodide, CoI

> Discuss briefly the difference between an orbit (as described by Bohr for hydrogen) and an orbital (as described by the more modern, wave mechanical picture of the atom).

> Why was Bohr’s theory for the hydrogen atom initially accepted, and why was it ultimately discarded?

> What are the essential points of Bohr’s theory of the structure of the hydrogen atom?

> When a tube containing hydrogen atoms is energized by passing several thousand volts of electricity into the tube, the hydrogen emits light that, when passed through a prism, resolves into the “bright line” spectrum sh

> What questions were left unanswered by Rutherford’s experiments?

> The energy levels of hydrogen (and other atoms) are said to be , which means that only certain energy values are allowed.

> How does the energy possessed by an emitted photon compare to the difference in energy levels that gave rise to the emission of the photon?

> Because a given element’s atoms emit only certain photons of light, only certain are occurring in those particular atoms.

> When an atom energy from outside, the atom goes from a lower energy state to a higher energy state.

> Calculate the percent by mass of the element listed first in the formulas for each of the following compounds. a. methane, CH4 b. sodium nitrate, NaNO3 c. carbon monoxide, CO d. nitrogen dioxide, NO2 e. 1-octanol, C8H18O f. calcium phosphate, Ca3(P

> How is the energy carried per photon of light related to the wavelength of the light? Does short-wavelength light carry more energy or less energy than long-wavelength light?

> Three elements have the electron configurations 1s22s22p63s2, 1s22s22p63s23p4, and 1s22s22p63s23p64s2. The first ionization energies of these elements (not in the same order) are 0.590, 0.999, and 0.738 MJ/mol. The atomic radii are 104, 160, and 197 pm.

> Compare the ionization energies of each pair of atoms. State the atom with the larger ionization energy for each pair. Symbol of Atom with the Larger lonization Energy Pair He and Kr Na and Al Cl and I

> Compare the atomic sizes of each pair of atoms. State the larger atom for each pair. Pair Symbol for Larger Atom F and B C and N B and Al

> Give the electron configurations for the following atoms. Use the noble gas notation. Element Electron Configuration K Ве Zr Se

> Identify the following three elements. a. The ground-state electron configuration is [Kr]5s24d105p4. b. The ground-state electron configuration is [Ar]4s23d104p2. c. An excited state of this element has the electron configuration 1s22s22p43s1.

> Give the electron configurations for the following atoms. Do not use the noble gas notation. Write out the complete electron configuration. Element Electron Configuration Ca Са B H Ве

> Which of the following statements is(are) true? a. The 2s orbital in the hydrogen atom is larger than the 3s orbital also in the hydrogen atom. b. The Bohr model of the hydrogen atom has been found to be incorrect. c. The hydrogen atom has quantized e

> An excited atom can release some or all of its excess energy by emitting a(n) and thus move to a lower energy state.

> Determine the maximum number of electrons that can have each of the following designations: 2f, 2dxy, 3p, 5dyz, and 4p.

> Calculate the percent by mass of each element in the following compounds. a. ZnO b. Na2S c. Mg(OH)2 d. H2O2 e. CaH2 f. K2O

> In each of the following sets of elements, indicate which element has the smallest atomic size. a. Ba, Ca, Ra b. P, Si, Al c. Rb, Cs, K

> In each of the following sets of elements, indicate which element shows the most active chemical behavior. a. B, Al, In b. Na, Al, S c. B, C, F

> Which of the following statements about the periodic table is false? a. Elements in the same column have similar reactivities because their valence electrons tend to be located in the same types of orbitals. b. A series of ions that are isoelectronic (

> Write the shorthand valence shell electron configuration of each of the following elements, basing your answer on the element’s location on the periodic table. a. nickel, Z = 28 b. niobium, Z = 41 c. hafnium, Z = 72 d. astatine, Z = 85

> Arrange the following atoms in order of increasing size (assuming all atoms are in their ground states). a. [Kr]5s24d105p6 b. [Kr]5s24d105p1 c. [Kr]5s24d105p3

> Using the symbol of the previous noble gas to indicate core electrons, write the valence shell electron configuration for each of the following elements. a. titanium, Z = 22 b. selenium, Z = 34 c. antimony, Z = 51 d. strontium, Z = 38

> Rank the following elements in order of increasing atomic size: Ge, S, F, Rb, Mn.

> What name is given to the series of ten elements in which the electrons are filling the 3d sublevel?

> How many valence electrons does each of the following atoms have? a. nitrogen, Z = 7 b. chlorine, Z = 17 c. sodium, Z = 11 d. aluminum, Z = 13

> What does the ground state of an atom represent?

> Calculate the percent by mass of each element in the following compounds. a. HClO3 b. UF4 c. CaH2 d. Ag2S e. NaHSO3 f. MnO2

> Write the complete orbital diagram for each of the following elements, using boxes to represent orbitals and arrows to represent electrons. a. scandium, Z = 21 b. sulfur, Z = 16 c. potassium, Z = 19 d. nitrogen, Z = 7

> Write the full electron configuration (1s22s2, etc.) for each of the following elements. a. bromine, Z = 35 b. xenon, Z = 54 c. barium, Z = 56 d. selenium, Z = 34

> Element X, which has a valence shell configuration of ns2np4, was isolated in a laboratory. Which of the following statements is(are) true concerning element X? a. Element X has chemical properties similar to those of the halogens. b. Element X has six

> Based on the ground-state electron configuration of iodine, how many electrons occupy the p and d orbitals?

> How does the attractive force that the nucleus exerts on an electron change with the principal energy level of the electron?

> Light waves move through space at a speed of meters per second.

> The energy of a photon of visible light emitted by an excited atom is the energy change that takes place within the atom itself.

> An atom has a small charged core called the nucleus, with charged electrons moving in the space around the nucleus.

> The random motions of the components of an object constitute the of that object.

> Consider the reaction B2H6(g) + 3O2(g) B2O3(s) + 3H2O(g) ∆H = -2035 kJ Calculate the amount of heat released when 54.0 g of diborane is combusted.

> Using the average atomic masses given inside the front cover of this book, calculate the mass in grams of each of the following samples. a. 2.17 moles of germanium, Ge b. 4.24 mmol of lead(II) chloride (1 mmol = 1/1000 mol) c. 0.0971 mole of ammonia,

> The specific heat capacity of graphite is 0.71 J/g °C. Calculate the energy required to raise the temperature of 2.4 moles of graphite by 25.0 °C.

> Which of the following processes is(are) exothermic? a. combustion of methane (e.g., Bunsen burner) b. melting of ice c. evaporation of acetone (e.g., fingernail polish remover) d. steam condensing on a cold surface

> Calculate the enthalpy change when 5.00 g of propane is burned with excess oxygen according to the reaction C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l) ∆H = -2221 kJ/mol

> Calculate ∆E for each of the following. a. q = -47 kJ, w = +88 kJ b. q = +82 kJ, w = +47 kJ c. q = +47 kJ, w = 0 d. In which of these cases do the surroundings do work on the system?

> A system releases 213 kJ of heat and has a calculated ∆E of -45 kJ. How much work was done on the system?

> If 7.24 kJ of heat is applied to a 952-g block of metal, the temperature increases by 10.7 °C. Calculate the specific heat capacity of the metal in J/g °C.

> If 10. J of heat is applied to 5.0-g samples of each of the substances listed in Table 10.1, which substance’s temperature will increase the most? Which substance’s temperature will increase the least?

> Calculate the amount of energy required (in joules) to heat 2.5 kg of water from 18.5 °C to 55.0 °C.

> The specific heat capacity of gold is 0.13 J/g °C. Calculate the specific heat capacity of gold in cal/g °C.

> What quantity of heat energy would have to be applied to a 25.1-g block of iron in order to raise the temperature of the iron sample by 17.5 °C? (See Table 10.1.) Table 10.1 The Specific Heat Capacities of Some Common Substances Specific