Solution: In each of the following sets, which atom or ion

We now have evidence that electron energy levels in the atoms are quantized. Some of this evidence is discussed in this chapter. What if energy levels in atoms were not quantized? What are some differences we would notice?

You have learned that each orbital is allowed two electrons, and this pattern is evident on the periodic table. What if each orbital was allowed three electrons? How would this change the appearance of the periodic table? For example, what would be the atomic numbers of the noble gases?

What if Bohrs model was correct? How would this affect the radial probability profiles as shown in Figs. 12.33 and 12.34?

Explain what it means for something to have wavelike properties; for something to have particulate properties. Electromagnetic radiation can be discussed in terms of both particles and waves. Explain the experimental verification for each of these views.

Defend and criticize Bohrs model. Why was it reasonable that such a model was proposed, and what evidence was there that it works? Why do we no longer believe in it?

The first ionization energy for magnesium is 735 kJ/mol. Which electron is this for? Estimate Zeff for this electron, and explain your reasoning. Calculate Zeff for this electron, and compare it to your estimate

The first four ionization energies for elements X and Y are shown below. The units are not kJ/mol. X Y First 170 200 Second 350 400 Third 1800 3500 Fourth 2500 5000 Identify the elements X and Y. There may be more than one answer, so explain completely.

Compare the first ionization energy of helium with its second ionization energy, remembering that both electrons come from the 1s orbital. Explain the difference without using actual numbers from the text.

Which has a larger second ionization energy, lithium or beryllium? Why?

Explain why a graph of ionization energy versus atomic number (across a row) is not linear. Where are the exceptions? Explain why they occur.

Without referring to your text, predict the trend of second ionization energies for the elements sodium through argon. Compare your answer with the data in Table 12.6. Explain any differences.

Account for the fact that the line that separates the metals from the nonmetals on the periodic table is diagonal downward to the right instead of horizontal or vertical

Explain the term electron from a quantum mechanical perspective, including a discussion of atomic radii, probabilities, and orbitals.

Choose the best response for the following. The ionization energy for the chlorine atom is equal in magnitude to the electron affinity for a. the Cl atom d. the F atom b. the Cl2 ion e. none of these c. the Cl1 ion Explain

Consider the following statement: The ionization energy for the potassium atom is negative because when K loses an electron to become K1 it achieves a noble gas electron configuration. Indicate what is incorrect. Explain.

What is the difference between Zeff and Z? When are they the same? Explain.

In going across a row of the periodic table, electrons are added and ionization energy generally increases. In going down a column of the periodic table, electrons are also being added but ionization energy generally decreases. Explain.

Explain the difference between the probability density distribution for an orbital and its radial probability.

How does the energy of a hydrogen 1s orbital compare with that of a lithium 1s orbital? Why? What is meant by the term energy of the orbital? What is its sign? Why? What is meant by the term lower in energy?

Which is larger, the hydrogen 1s orbital or the lithium 1s orbital? Why? Which has the larger radius, the hydrogen atom or the lithium atom? Why?

Is the following statement true or false: The hydrogen atom has a 3s orbital. Explain

Which is higher in energy: the 2s or 2p orbital in hydrogen? Is this also true for helium? Explain

Prove mathematically that it is more energetically favorable for a fluorine atom to take an electron from a sodium atom than for a fluorine atom to take an electron from another fluorine atom.

Microwave radiation has a wavelength on the order of 1.0 cm. Calculate the frequency and the energy of a single photon of this radiation. Calculate the energy of an Avogadros number of photons (called an einstein) of this electromagnetic radiation.

Consider the following waves representing electromagnetic radiation: 1.6 103 m Wave a Wave b Which wave has the longer wavelength? Calculate the wavelength. Which wave has the higher frequency and larger photon energy? Calculate these values. Which wave has the greater velocity? What type of electromagnetic radiation does each wave represent?

Octyl methoxycinnamate and oxybenzone are common ingredients in sunscreen applications. These compounds work by absorbing ultraviolet (UV) B light (wavelength 280320 nm), the UV light most associated with sunburn symptoms. What frequency range of light do these compounds absorb?

Human color vision is produced by the nervous system based on how three different cone receptors interact with photons of light in the eye. These three different types of cones interact with photons of different frequency light, as indicated in the following table: Cone Type Range of Light Frequency Detected S 6.007.49 3 1014 s21 M 4.766.62 3 1014 s21 L 4.286.00 3 1014 s21 What wavelength ranges (and corresponding colors) do the three types of cones detect?

One type of electromagnetic radiation has a frequency of 107.1 MHz, another type has a wavelength of 2.12 3 10210 m, and another type of electromagnetic radiation has photons with energy equal to 3.97 3 10219 J/photon. Identify each type of electromagnetic radiation, and place them in order of increasing photon energy and increasing frequency

Carbon absorbs energy at a wavelength of 150. nm. The total amount of energy emitted by a carbon sample is 1.98 3 105 J. Calculate the number of carbon atoms present in the sample, assuming that each atom emits one photon

A carbonoxygen double bond in a certain organic molecule absorbs radiation that has a frequency of 6.0 3 1013 s21. a. What is the wavelength of this radiation? b. To what region of the spectrum does this radiation belong?c. What is the energy of this radiation per photon? Per mole of photons? d. A carbonoxygen bond in a different molecule absorbs radiation with frequency equal to 5.4 3 1013 s21. Is this radiation more or less energetic?

X rays have wavelengths on the order of 1 3 10210 m. Calculate the energy of 1.0 3 10210 m X rays in units of kilojoules per mole of X rays. AM radio waves have wavelengths on the order of 1 3 104 m. Calculate the energy of 1.0 3 104 m radio waves in units of kilojoules per mole of radio waves. Consider that the bond energy of a carboncarbon single bond found in organic compounds is 347 kJ/mol. Would X rays and/or radio waves be able to disrupt organic compounds by breaking carboncarbon single bonds?

The work function of an element is the energy required to remove an electron from the surface of the solid. The work function for lithium is 279.7 kJ/mol (that is, it takes 279.7 kJ of energy to remove 1 mole of electrons from 1 mole of Li atoms on the surface of Li metal). What is the maximum wavelength of light that can remove an electron from an atom in lithium metal?

Ionization energy is the energy required to remove an electron from an atom in the gas phase. The ionization energy of gold is 890.1 kJ/mol. Is light with a wavelength of 225 nm capable of ionizing a gold atom (removing an electron) in the gas phase?

It takes 208.4 kJ of energy to remove 1 mole of electrons from the atoms on the surface of rubidium metal. If rubidium metal is irradiated with 254-nm light, what is the maximum kinetic energy the released electrons can have?

What experimental evidence supports the quantum theory of light? Explain the waveparticle duality of all matter. For what size particle must one consider both the wave and the particle properties?

Explain the photoelectric effect

Calculate the de Broglie wavelength for each of the following. a. an electron with a velocity 10.% of the speed of light b. a tennis ball (55 g) served at 35 m/s (,80 mi/h)

Neutron diffraction is used in determining the structures of molecules. a. Calculate the de Broglie wavelength of a neutron moving at 1.00% of the speed of light. b. Calculate the velocity of a neutron with a wavelength of 75 pm (1 pm 5 10212 m).

Calculate the velocities of electrons with de Broglie wavelengths of 1.0 3 102 nm and 1.0 nm, respectively.

An atom of a particular element is traveling at 1% of the speed of light. The de Broglie wavelength is found to be 3.31 3 1023 pm. Which element is this?

Characterize the Bohr model of the atom. In the Bohr model, what do we mean when we say something is quantized? How does the Bohr model of the hydrogen atom explain the hydrogen emission spectrum? Why is the Bohr model fundamentally incorrect?

The following is an energy-level diagram illustrating three different electronic transitions in the Bohr hydrogen atom. E 4 5 3 2 1 n a. Explain why the energy levels get closer together as they increase. Provide mathematical support for this. b. Verify that the colors given in the diagram are correct. Provide mathematical support.

Consider only the transitions involving the first four energy levels for a hydrogen atom: n = 4 n = 3 n = 2 n = 1 a. How many emissions are possible for an electron in the n 5 4 level as it goes to the ground state? b. Which electronic transition is the lowest energy? c. Which electronic transition corresponds to the shortest wavelength emission?

Calculate the longest and shortest wavelengths of light emitted by electrons in the hydrogen atom that begin in the n 5 6 state and then fall to states with smaller values of n.

Calculate the wavelength of light emitted when each of the following transitions occur in the hydrogen atom. What type of electromagnetic radiation is emitted in each transition? a. n 5 4 n n 5 3 b. n 5 5 n n 5 4 c. n 5 5 n n 5 3

Assume that a hydrogen atoms electron has been excited to the n 5 5 level. How many different wavelengths of light can be emitted as this excited atom loses energy?

What is the maximum wavelength of light capable of removing an electron from a hydrogen atom in the energy states characterized by n 5 1 and n 5 3?

An electron is excited from the ground state to the n 5 3 state in a hydrogen atom. Which of the following statements are true? Correct any false statements. a. It takes more energy to ionize (remove) the electron from n 5 3 than from the ground state. b. The electron is farther from the nucleus on average in the n 5 3 state than in the ground state. c. The wavelength of light emitted if the electron drops from n 5 3 to n 5 2 is shorter than the wavelength of light emitted if the electron falls from n 5 3 to n 5 1. d. The wavelength of light emitted when the electron returns to the ground state from n 5 3 is the same as the wavelength of light absorbed to go from n 5 1 to n 5 3. e. The first excited state corresponds to n 5 3.

Does a photon of visible light (l 5 400700 nm) have sufficient energy to excite an electron in a hydrogen atom from the n 5 1 to the n 5 5 energy state? From the n 5 2 to the n 5 6 energy state?

An excited hydrogen atom emits light with a wavelength of 397.2 nm to reach the energy level for which n 5 2. In which principal quantum level did the electron begin?

An excited hydrogen atom with an electron in the n 5 5 state emits light having a frequency of 6.90 3 1014 s21. Determine the principal quantum level for the final state in this electronic transition.

Consider an electron for a hydrogen atom in an excited state. The maximum wavelength of electromagnetic radiation that can completely remove (ionize) the electron from the H atom is 1460 nm. Determine the initial excited state for the electron (n 5 ?).

Calculate the energy (in kJ/mol) required to remove the electron in the ground state for each of the following one-electron species using the Bohr model. a. H b. He1 c. Li21 d. C51 e. Fe251

One of the emission spectral lines for Be31 has a wavelength of 253.4 nm for an electronic transition that begins in the state with n 5 5. What is the principal quantum number of the lower-energy state corresponding to this emission?

The Heisenberg uncertainty principle can be expressed in the form DE # Dt $ U 2 where E represents energy and t represents time. Show that the units for this form are the same as the units for the form used in this chapter: Dx # Dp $ U 2

Using the Heisenberg uncertainty principle, calculate Dx for each of the following. a. an electron with Dv 5 0.100 m/s b. a baseball (mass 5 145 g) with Dv 5 0.100 m/s How does the answer in part a compare with the size of a hydrogen atom? How does the answer in part b correspond to the size of a baseball?

We can represent both probability and radial probability versus distance from the nucleus for a hydrogen 1s orbital as depicted below. Probability ( R2) Distance from nucleus (r) Radial probability (4r2R2 ) Distance from nucleus (r) What does each graph tell us about the electron in a hydrogen 1s orbital? Describe the significance of the radial probability distribution

Discuss why a function of the type A cos(Lx) is not an appropriate solution for the particle in a one-dimensional box

Calculate the wavelength of the electromagnetic radiation required to excite an electron from the ground state to the level with n 5 5 in a one-dimensional box 40.0 pm in length.

An electron in a one-dimensional box requires a wavelength of 8080 nm to excite an electron from the n 5 2 to the n 5 3 energy level. Calculate the length of this box

An electron in a 10.0-nm one-dimensional box is excited from the ground state into a higher-energy state by absorbing a photon of electromagnetic radiation with a wavelength of 1.374 3 1025 m. Determine the final energy state for this transition.

Discuss what happens to the energy levels for an electron trapped in a one-dimensional box as the length of the box increases.

What is the total probability of finding a particle in a one-dimensional box in level n 5 3 between x 5 0 and x 5 L/6?

Which has the lowest (ground-state) energy, an electron trapped in a one-dimensional box of length 1026 m or one with length 10210 m?

What are quantum numbers? What information do we get from the quantum numbers n, ,, and m,? We define a spin quantum number (ms), but do we know that an electron literally spins?

How do 2p orbitals differ from each other? How do 2p and 3p orbitals differ from each other? What is a nodal surface in an atomic orbital?

Identify each of the following orbitals, and determine the n and l quantum numbers. Explain your answers. a. b. Node c. y z x +

Which of the following orbital designations are incorrect: 1s, 1p, 7d, 9s, 3f, 4f, 2d?

Which of the following sets of quantum numbers are not allowed? For each incorrect set, state why it is incorrect. a. n 5 3, , 5 3, m, 5 0, ms 5 21 2 b. n 5 4, , 5 3, m, 5 2, ms 5 21 2 c. n 5 4, , 5 1, m, 5 1, ms 5 11 2 d. n 5 2, , 5 1, m, 5 21, ms 5 21 e. n 5 5, , 5 24, m, 5 2, ms 5 11 2 f. n 5 3, , 5 1, m, 5 2, ms 5 21 2

The following sets of quantum numbers are not correct. Which quantum number is not consistent with the others, or is just wrong? a. m, 5 23, n 5 3, ms 5 21 2, , 5 1 b. ms 5 11 2, n 5 3, , 5 1, m, 5 2

How many orbitals can have the designation 5p, 3dz2, 4d, n 5 5, and n 5 4?

How many electrons in an atom can have the designation 1p, 6dx22y2, 4f, 7py, 2s, and n 5 3?

What is the physical significance of the value of c2 at a particular point in an atomic orbital?

In defining the sizes of orbitals, why must we use an arbitrary value, such as 90% of the probability of finding an electron in that region?

From the diagrams of 2p and 3p orbitals in Fig. 12.19 and Fig. 12.20, respectively, draw a rough graph of the square of the wave function for these orbitals in the direction of one of the lobes

The wave function for the 2pz orbital in the hydrogen atom is c2pz 5 1 4!2pa Z a0 b 3/2 se2s/2 cos u where a0 is the value for the radius of the first Bohr orbit in meters (5.29 3 10211), s is Zr/a0, r is the value for the distance from the nucleus in meters, and u is an angle. Calculate the value of c2pz 2 at r 5 a0 for u 5 0 (z axis) and for u 5 908 (xy plane).

For hydrogen atoms, the wave function for the state n 5 3, , 5 0, and m, 5 0 is c300 5 1 81!3pa 1 a0 b 3/2 127 2 18s 1 2s22e2s/3 where s 5 r/a0 and a0 is the Bohr radius (5.29 3 10211 m). Calculate the position of the nodes for this wave function

Total radial probability distributions for the helium, neon, and argon atoms are shown in the following graph. How can the shapes of these curves be interpreted in terms of electron configurations, quantum numbers, and nuclear charges? Radial electron density 0 Distance from nucleus ()

The relative orbital levels for the hydrogen atom can be represented as E 3s 2s 1s 3p 2p 3d Draw the relative orbital energy levels for atoms with more than one electron, and explain your answer. Also explain how the following radial probability distributions support your answer. Distance from the nucleus Radial probability 2p 2s Radial probability Distance from the nucleus 3s 3p 3d Penetration 77. What is the difference between core electrons and valence electrons? Why do we emphasize the valence electrons in an atom when discussing atomic properties? What is the relationship between valence electrons and elements in the same group of the periodic table?

The periodic table consists of four blocks of elements that correspond to s, p, d, and f orbitals being filled. After f orbitals come g and h orbitals. In theory, if a g block and an h block of elements existed, how long would the rows of g and h elements be in this theoretical periodic table?

What is the maximum number of electrons in an atom that can have these quantum numbers? a. n 5 4 b. n 5 5, m, 5 11 c. n 5 5, ms 5 11 2 d. n 5 3, , 5 2 e. n 5 2, , 5 1 f. n 5 0, , 5 0, m, 5 0 g. n 5 2, , 5 1, m, 5 21, ms 5 21 2 h. n 5 3, ms 5 11 2 i. n 5 2, , 5 2 j. n 5 1, , 5 0, m, 5 0

The elements of Si, Ga, As, Ge, Al, Cd, S, and Se are all used in the manufacture of various semiconductor devices. Write the expected electron configurations for these atoms

Write the expected electron configurations for the following atoms: Sc, Fe, P, Cs, Eu, Pt, Xe, and Br.

Write the expected electron configurations for each of the following atoms: Cl, As, Sr, W, Pb, and Cf

Using Fig. 12.29, list elements (ignore the lanthanides and actinides) that have ground-state electron configurations that differ from those we would expect from their positions in the periodic table.

Write the expected ground-state electron configuration for the following. a. the element with one unpaired 5p electron that forms a covalent compound with fluorine b. the (as yet undiscovered) alkaline earth metal after radium c. the noble gas with electrons occupying 4f orbitals d. the first-row transition metal with the most unpaired electrons

For elements 136, there are two exceptions to the filling order as predicted from the periodic table. Draw the atomic orbital diagrams for the two exceptions, and indicate how many unpaired electrons are present.

Given the valence electron orbital level diagram and the description, identify the element or ion. a. A ground state atom 3s 3p b. An atom in an excited state (assume two electrons occupy the 1s orbital) 2s 2p c. A ground state ion with a charge of 21 4s 4p

How many valence electrons do each of the following elements have, and what are the specific valence electrons for each element? a. Ca d. In b. O e. Ar c. element 117 f. Bi

In the ground state of mercury (Hg), a. how many electrons occupy atomic orbitals with n 5 3? b. how many electrons occupy d atomic orbitals? c. how many electrons occupy pz atomic orbitals? d. how many electrons have spin up (ms 5 11 2)?

In the ground state of element 115, Uup, a. how many electrons have n 5 5 as one of their quantum numbers? b. how many electrons have , 5 3 as one of their quantum numbers? c. how many electrons have m, 5 1 as one of their quantum numbers? d. how many electrons have ms 5 21 2 as one of their quantum numbers?

Give possible values for the quantum numbers of the valence electrons in an atom of titanium (Ti).

Give a possible set of values of the four quantum numbers for all the electrons in a boron atom and a nitrogen atom if each is in the ground state

How many unpaired electrons are present in each of the first-row transition metals in the ground state?

One bit of evidence that the quantum mechanical model is correct lies in the magnetic properties of matter. Atoms with unpaired electrons are attracted by magnetic fields and thus are said to exhibit paramagnetism. The degree to which this effect is observed is directly related to the number of unpaired electrons present in the atom. Consider the ground-state electron configurations for Li N, Ni, Te, Ba, and Hg. Which of these atoms would be expected to be paramagnetic, and how many unpaired electrons are present in each paramagnetic atom?

Which of elements 136 have two unpaired electrons in the ground state?

Which of elements 136 have one unpaired electron in the ground state?

A certain oxygen atom has the electron configuration 1s22s22px 22py 2. How many unpaired electrons are present? Is this an excited state for oxygen? In going from this state to the ground state, would energy be released or absorbed?

How many unpaired electrons are present in each of the following in the ground state: O, O1, O2, Os, Zr, S, F, Ar?

Which of the following electron configurations correspond to an excited state? Identify the atoms, and write the ground-state electron configuration where appropriate. a. 1s22s23p1 c. 1s22s22p43s1 b. 1s22s22p6 d. [Ar]4s23d54p1

Using the element phosphorus as an example, write equations for the processes in which the energy change will correspond to the ionization energy and to the electron affinity

Explain why the first ionization energy tends to increase as one proceeds from left to right across a period. Why is the first ionization energy of aluminum lower than that of magnesium and the first ionization energy of sulfur lower than that of phosphorus?

Why do the successive ionization energies of an atom always increase? Note the successive ionization energies for silicon given in Table 12.6. Would you expect to see any large jumps between successive ionization energies of silicon as you removed all the electrons, one by one, beyond those shown in the table?

The radius trend and the ionization energy trend are exact opposites. Does this make sense? Define electron affinity. Electron affinity values are both exothermic (negative) and endothermic (positive). However, ionization energy values are always endothermic (positive). Explain.

Arrange the following groups of atoms in order of increasing size. a. Te, S, Se d. Rb, Na, Be b. K, Br, Ni e. Sr, Se, Ne c. Ba, Si, F f. Fe, P, O

Arrange the atoms in Exercise 103 in order of increasing first ionization energy.

In each of the following sets, which atom or ion has the smallest ionization energy? a. Ca, Sr, Ba d. S22, S, S21 b. K, Mn, Ga e. Cs, Ge, Ar c. N, O, F

In each of the following sets, which atom or ion has the smallest radius? a. H, He b. Cl, In, Se c. element 120, element 119, element 117 d. Nb, Zn, Si e. Na2, Na, Na1

The first ionization energies of As and Se are 0.947 MJ/mol and 0.941 MJ/mol, respectively. Rationalize these values in terms of electron configurations.

Rank the elements Be, B, C, N, and O in order of increasing first ionization energy. Explain your reasoning

We expect the atomic radius to increase down a group in the periodic table. Can you suggest why the atomic radius of hafnium breaks this rule? (See the following data.) Element Atomic Radius () Element Atomic Radius () Sc 1.57 Ti 1.477 Y 1.693 Zr 1.593 La 1.915 Hf 1.47

Answer the following questions based on the given electron configurations, and identify the elements. a. Arrange these atoms in order of increasing size: [Kr]5s24d105p6; [Kr]5s24d105p1; [Kr]5s24d105p3. b. Arrange these atoms in order of decreasing first ionization energy: [Ne]3s23p5; [Ar]4s23d104p3; [Ar]4s23d104p5.

Predict some of the properties of element 117 (symbol Uus following conventions proposed by the International Union of Pure and Applied Chemistry [IUPAC]). a. What will be its electron configuration? b. What element will it most resemble chemically? c. What will be the formulas of the neutral binary compounds it forms with sodium, magnesium, carbon, and oxygen? d. What oxyanions would you expect Uus to form?

Consider the following ionization energies for aluminum. Al(g) 88n Al1(g) 1 e2 I1 5 580 kJ/mol Al1(g) 88n Al21(g) 1 e2 I2 5 1815 kJ/mol Al21(g) 88n Al31(g) 1 e2 I3 5 2740 kJ/mol Al31(g) 88n Al41(g) 1 e2 I4 5 11,600 kJ/mol a. Account for the increasing trend in the values of the ionization energies. b. Explain the large increase between I3 and I4. c. Which one of the four ions has the greatest electron affinity? Explain. d. List the four aluminum ions given in the preceding reactions in order of increasing size, and explain your ordering. (Hint: Remember that most of the size of an atom or ion is due to its electrons.)

The following graph plots the first, second, and third ionization energies for Mg, Al, and Si. Number of electrons removed 1 2 3 Ionization energy (kJ/mol) Without referencing the text, which plot corresponds to which element? In one of the plots, there is a huge jump in energy between I2 and I3, unlike in the other two plots. Explain this phenomenon.

Order each of the following sets from the least exothermic electron affinity to the most. a. F, Cl, Br, I b. N, O, F

In the second row of the periodic table, Be, N, and Ne all have endothermic (unfavorable) electron affinities, whereas the other second-row elements have exothermic (favorable) electron affinities. Rationalize why Be, N, and Ne have unfavorable electron affinities

Which has the more negative electron affinity, the oxygen atom or the O2 ion? Explain your answer.

The electron affinity for sulfur is more exothermic than that for oxygen. How do you account for this?

The electron affinities of the elements from aluminum to chlorine are 244 kJ/mol, 2120 kJ/mol, 274 kJ/mol, 2200.4 kJ/mol, and 2348.7 kJ/mol, respectively. Rationalize the trend in these values.

Use data in this chapter to determine the following. a. the electron affinity of Mg21 b. the electron affinity of Al1 c. the ionization energy of Cl2 d. the ionization energy of Cl e. the electron affinity of Cl1

For each of the following pairs of elements, (C and N) (Ar and Br) (Mg and K) (F and Cl) pick the one with a. the more favorable (exothermic) electron affinity b. the higher ionization energy c. the larger size

Does the information on alkali metals in Table 12.9 of the text confirm the general periodic trends in ionization energy and atomic radius? Explain.

An ionic compound of potassium and oxygen has the empirical formula KO. Would you expect this compound to be potassium(II) oxide or potassium peroxide? Explain

Complete and balance the equations for the following reactions. a. Li(s) 1 N2(g) 88n c. Cs(s) 1 H2O(l) 88n b. Rb(s) 1 S(s) 88n d. Na(s) 1 Cl2(g) 88n

Cesium was discovered in natural mineral waters in 1860 by R. W. Bunsen and G. R. Kirchhoff, using the spectroscope they invented in 1859. The name comes from the Latin word caesius, meaning sky blue, which describes the prominent blue line observed for this element at 455.5 nm. Calculate the frequency and energy of a photon of this light.

The bright yellow light emitted by a sodium vapor lamp consists of two emission lines at 589.0 nm and 589.6 nm. What are the frequency and the energy of a photon of light at each of these wavelengths? What are the energies in kJ/mol?

Give the name and formula of the binary compound formed by each of the following pairs of elements. a. Li and N d. Li and P b. Na and Br e. Rb and H c. K and S f. Na and H

Predict the atomic number of the next alkali metal after francium, and give its ground-state electron configuration.

Spectroscopists use emission spectra to confirm the presence of an element in materials of unknown composition. How is this possible?

Without looking at data in the text, sketch a qualitative graph of the third ionization energy versus atomic number for the elements Na through Ar, and explain your graph.