3 questions, one for each idea where we can. Answer them, then see which ideas to fix.
Question 1 of 3
The table gives the ground-state electron configurations of atoms of five elements labeled Q, R, T, X and Z. Which two of the elements are in the same group of the periodic table?
Answer and reasoning
AQ and T A student who places every atom ending in two s electrons in group 2 picks this. Q (1s², helium) has a completely filled shell and is a noble gas; T (beryllium) has two valence electrons outside a filled 1s core.
BR and T A student who thinks elements with the same number of occupied shells are in the same group picks this. R and T both have two occupied shells, so they are in the same period, period 2, but they have one and two valence electrons respectively.
CR and XCorrect R (1s² 2s¹, lithium) and X (1s² 2s² 2p⁶ 3s¹, sodium) each have one valence electron in an s subshell outside a filled-shell core, so both are in group 1.
DX and Z A student who counts only the electrons in the last-filled subshell as valence electrons picks this, because X ends in 3s¹ and Z in 2p¹. Z (boron) has three valence electrons, 2s² 2p¹, so it is in group 13, while X is in group 1.
Which statement describes the process whose energy change is the first ionization energy of magnesium?
Answer and reasoning
AAn atom in solid Mg loses two electrons, forming an aqueous Mg²⁺ ion A student who thinks ionization means forming the ion the element normally has picks this. The first ionization energy refers to removing one electron from a gaseous atom to form a gaseous 1+ ion.
BAn atom in solid Mg loses one electron, forming a gaseous Mg⁺ ion A student who thinks ionization energy refers to the element in its usual state picks this. Starting from the solid would add the energy needed to separate a Mg atom from the metal; ionization energy is defined for gaseous atoms.
CA gaseous Mg atom takes in one electron, forming a gaseous Mg⁻ ion A student who treats ionization energy and electron affinity as the same process picks this. Adding an electron to a gaseous atom is the process of electron affinity; ionization removes an electron.
DA gaseous Mg atom loses one electron, forming a gaseous Mg⁺ ionCorrect The first ionization energy is the energy needed to remove one electron from an isolated gaseous atom in its ground state, forming a gaseous 1+ ion.
Working First ionization: remove one electron from one gaseous atom in its ground state: Mg(g) → Mg⁺(g) + e⁻. Charge balances (0 = +1 − 1) and the states are gaseous.
Astatine, At, lies directly below iodine in group 17. All of its isotopes are highly radioactive and only tiny amounts can be made, so few of its properties have been measured. A team plans to measure one atomic property of astatine. Based on periodic trends, which result should the team predict?
Answer and reasoning
AIts first ionization energy will be higher than that of iodine. A student who thinks more protons always mean a stronger pull on the outer electrons picks this. Astatine's extra protons are outweighed by the greater distance and shielding of its valence electrons.
BIts first ionization energy will be equal to the value for iodine. A student who thinks elements in the same group have the same values picks this. Within a group, first ionization energy decreases from top to bottom.
CIts first ionization energy will be lower than that of iodine.Correct Astatine's valence electrons are in the sixth shell, farther from the nucleus and more shielded than iodine's fifth-shell electrons, so less energy is needed to remove one. (The value measured in 2013, about 899 kJ/mol, is below iodine's 1008 kJ/mol.)
DIts atoms will be smaller than the atoms of iodine are. A student who thinks a heavier nucleus pulls the electrons in more strongly picks this. The attraction depends on charge and distance, not mass, and astatine's valence electrons occupy a higher shell, so its atoms are larger than iodine's.
In preparation: 0 of 3 sections compiled and reviewed. The rest show key terms and common mistakes from our question bank until they are.
1.7.A.1 Periodicity Fix
Periodicity
The recurring pattern of properties seen when the elements are arranged in order of atomic number. Elements in the same group (column) generally have the same arrangement of valence electrons, in successively higher shells (helium, 1s², is the exception in group 18), so their properties recur from period to period.
Valence electrons
The electrons in the outermost occupied shell of a main-group atom (its highest-numbered s and p subshells). Atoms of main-group elements in the same group have the same number of valence electrons (helium, with two, is the exception in group 18); the electrons in completely filled inner shells are core electrons.
Students often think Elements in the same period have similar properties, because their atoms have the same number of occupied shells. In fact No. Elements in the same period have the same number of occupied shells, but different numbers of valence electrons, and their properties change across the period. Elements with similar properties are in the same group (column), where atoms have the same arrangement of valence electrons.
Students often think Elements whose atoms end in two s electrons all belong together, so helium (1s²) has the properties of the group 2 elements. In fact No. Helium's two electrons fill its first shell completely, so helium, like neon and argon, has a completely filled valence shell and is unreactive. The group 2 elements have two valence electrons outside a filled shell, which they lose readily.
1.7.A.2 Coulomb's law (in atoms) Fix
Coulomb's law (in atoms)
The force between two charges is proportional to the product of the charges divided by the square of the distance between them, F ∝ q₁q₂/r². Applied to an atom, an electron is attracted more strongly when the positive charge acting on it is larger and when it is closer to the nucleus.
Shell model
A model of the atom in which electrons occupy shells at increasing average distances from the nucleus (n = 1, 2, 3 …). Electrons in a higher shell are, on average, farther from the nucleus.
Shielding
The reduction in the attraction an electron experiences from the nucleus caused by repulsion from other electrons, mainly those in shells closer to the nucleus. Electrons in the same shell shield one another only weakly.
Effective nuclear charge
The net positive charge actually experienced by an electron: the nuclear charge reduced by shielding from the other electrons. Across a period it increases, because protons are added while the added electrons go into the same shell and shield poorly.
First ionization energy
The energy required to remove the most loosely held electron from a gaseous atom in its ground state, X(g) → X⁺(g) + e⁻. It generally increases across a period and decreases down a group.
Successive ionization energies
The energies needed to remove a first, second, third … electron from the same atom, each from an ion that has already lost the earlier electrons. They increase steadily while valence electrons are removed and jump sharply when the first core electron is removed.
Atomic radius
A measure of the size of an atom. It generally decreases across a period, because the effective nuclear charge increases while the valence electrons stay in the same shell, and increases down a group, because the valence electrons occupy shells farther from the nucleus.
Ionic radius
A measure of the size of an ion. A cation is smaller than its parent atom, and an anion is larger than its parent atom; in a series of ions with the same number of electrons, the radius decreases as the nuclear charge increases.
Electron affinity
The energy change when an electron is added to a gaseous atom, X(g) + e⁻ → X⁻(g). For most nonmetal atoms energy is released, and more is released when the added electron is attracted more strongly to the nucleus.
Electronegativity
The ability of an atom in a chemical bond to attract the shared electrons toward itself. For the representative elements it increases across a period and decreases down a group; fluorine has the highest value.
Students often think The more protons an atom's nucleus has, the more strongly it attracts its outer electrons, whatever shell they occupy. In fact No. The attraction on an outer electron depends on the net charge it experiences after shielding and on its distance from the nucleus. Down a group the number of protons increases, but the outer electrons are in higher shells, farther away and shielded by more inner electrons, so they are generally held less strongly.
Students often think Coulombic attraction decreases in simple proportion to distance, so doubling the distance halves the force. In fact No. Coulomb's law has the distance squared in the denominator, so doubling the distance makes the force one-quarter as great, if the charges are unchanged.
1.7.A.3 Estimating properties from periodicity Fix
Estimating properties from periodicity
Using an element's position in the periodic table, and measured values for its neighbors in the same group and period, to predict or estimate a property that has not been measured.
13 more questions. Every wrong answer here is a real mistake students make, and you see why it is wrong as soon as you answer.
Question 1 of 13
An element has the ground-state electron configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 4s². The chemical properties of this element are most similar to those of which of the following elements?
Answer and reasoning
AHe A student who groups every atom that ends in two s electrons with group 2 picks this. Helium's 1s² completely fills its only shell, so it is an unreactive noble gas; the element shown has two valence electrons outside a filled core and readily loses them.
BBr A student who thinks elements in the same period behave alike picks this. Bromine is in period 4 like this element, but it has seven valence electrons, not two, and is a reactive nonmetal; similar properties go with the same group.
CMgCorrect The element (calcium) has two valence electrons, 4s², outside a filled-shell core. Magnesium, 1s² 2s² 2p⁶ 3s², has the same arrangement of valence electrons one shell lower, so the two are in group 2 and behave similarly, for example both forming 2+ ions.
DAr A student who thinks a filled subshell makes an atom stable and unreactive picks this, because 4s² is a filled subshell. The two 4s electrons are valence electrons outside an argon-like core, and they are lost readily; argon has a completely filled valence shell and no such electrons.
The graph shows the first ionization energies of the elements with atomic numbers 1 to 20. Which claim is best supported by the data and the shell model of the atom?
Answer and reasoning
AA new shell starts after every element whose value is higher than the next element's, because each drop shows a new shell. A student who thinks every drop marks a new shell picks this. The small dips after Z = 4, 7, 12 and 15 do not start new shells: the electrons removed from B and Al are the first in a p subshell, and those removed from O and S come from a p orbital that holds two electrons, all within an existing shell. Only the large drops after Z = 2, 10 and 18 mark new shells.
BThe values rise as Z increases from 1 to 20, because each added proton attracts the outer electrons more strongly. A student who thinks more protons always means a stronger pull on the outer electrons picks this. The graph contradicts it: potassium (Z = 19) has a far lower value than hydrogen (Z = 1), because its outer electron is in the fourth shell, far from the nucleus and heavily shielded.
CAtoms in the same group, such as Z = 2, 10 and 18, have about equal values, because they have similar configurations. A student who thinks elements in a group have the same values picks this. The noble gases have the highest value in each period, but the values fall down the group, from 2372 kJ/mol for He to 1521 kJ/mol for Ar, as the outer electrons are farther from the nucleus.
DThe atoms with Z = 3, 11 and 19 each start a new shell, because each has a far lower value than the element just before it.Correct Each of these elements follows a noble gas (Z = 2, 10, 18) and has a first ionization energy roughly a quarter or less of that noble gas's. Their outermost electron is alone in a new shell, farther from the nucleus and shielded by a complete inner shell, so it is removed far more easily; this is the repeating pattern the periodic table is built on.
In a simple model, the force of attraction between the nucleus and a valence electron is proportional to q/r², where q is the net positive charge acting on the electron and r is the electron's average distance from the nucleus. For the valence electron of hypothetical atom Y, q is twice as great and r is twice as great as for the valence electron of hypothetical atom X. If the force on the valence electron of X is F, what is the force on the valence electron of Y?
Answer and reasoning
A0.25 F A student who thinks only the distance of an electron from the nucleus decides the attraction picks this, applying the factor (1/2)² for the doubled distance but ignoring the doubled charge. Both factors apply: (2) × (1/2)² = 0.50.
B0.50 FCorrect Force ∝ q/r². Doubling q multiplies the force by 2, and doubling r multiplies it by (1/2)² = 1/4, so the force on Y's valence electron is 2 × 1/4 = 0.50 times that on X's.
C1.00 F A student who thinks the force falls in simple proportion to distance picks this: doubling q (× 2) and doubling r (÷ 2) would cancel. The distance is squared, so doubling r divides the force by 4, not 2.
D2.00 F A student who thinks the charge alone decides the attraction picks this, doubling the force for the doubled charge and ignoring the doubled distance. The doubled distance divides the force by 4.
Working F ∝ q/r². FY/FX = (qY/qX) × (rX/rY)² = 2 × (1/2)² = 0.50, so FY = 0.50 F.
A student wants to use Coulomb's law to decide whether a sodium atom or a potassium atom attracts its valence electron more strongly. Which two quantities should the student compare for the two atoms?
Answer and reasoning
AThe net charge acting on the valence electron and the electron's average distance from the nucleusCorrect Coulomb's law makes the attraction proportional to the charges and inversely proportional to the square of the distance between them. For a valence electron, the relevant charge is the net charge it experiences after shielding by inner electrons, and the relevant distance is its average distance from the nucleus.
BThe mass of each nucleus and the number of neutrons that each of the two nuclei contains A student who thinks a heavier nucleus pulls harder on its electrons picks this. The attraction is Coulombic and depends on charge, not mass; neutrons add mass but no charge.
CThe number of valence electrons and the number of electrons still needed to complete an octet A student who thinks an atom holds its electrons tightly because it is close to a full octet picks this. Coulomb's law involves charge and distance; sodium and potassium have the same number of valence electrons, and the difference between them comes from distance and shielding.
DThe number of protons in each nucleus and the number of electrons among which their attraction is shared A student who thinks a nucleus has a fixed amount of attraction shared out among its electrons picks this. Each electron is attracted according to the net charge it experiences and its distance from the nucleus; the attraction is not divided up among the electrons.
Working Coulomb's law: F ∝ q₁q₂/r². For a valence electron, q₁ is the net positive charge acting on it (the nuclear charge reduced by shielding from inner electrons), q₂ is the charge of the electron (the same in both atoms) and r is its average distance from the nucleus. The two quantities that differ between Na and K, and so must be compared, are the net charge and the distance.
The first ionization energy of nitrogen is 1402 kJ/mol and that of oxygen is 1314 kJ/mol. Which statement best explains the change in first ionization energy from nitrogen to oxygen?
Answer and reasoning
AIn O the electron removed shares a 2p orbital with a second electron, so repulsion makes it easier to remove.Correct N (2p³) has one electron in each of its three 2p orbitals. In O (2p⁴) one 2p orbital holds two electrons, and the repulsion between them means that less energy is needed to remove one of them, so the first ionization energy falls slightly from N to O even though O has one more proton.
BIn O the electron removed is in a new, higher-energy subshell, which is farther from the nucleus than the 2p subshell. A student who thinks every drop in ionization energy means a new shell or subshell picks this. The electrons removed from N and from O are both 2p electrons; a new subshell explains the earlier dip from Be to B, where the electron removed is the first 2p electron.
CO atoms have more electrons than N atoms, so the O atoms are larger and their outer electrons are farther away. A student who thinks more electrons make an atom larger picks this. O atoms are smaller than N atoms: the added electron goes into the same shell, and the added proton pulls all of the electrons closer.
DIn O the attraction of the nucleus is shared among more electrons than in N, so each is easier to remove. A student who thinks a nucleus has a fixed amount of attraction that is shared out among its electrons picks this. Each electron is attracted according to the net charge it experiences and its distance from the nucleus; O has one more proton than N, and the fall in ionization energy comes from repulsion between the two electrons that share one 2p orbital.
The table gives the first five ionization energies of an element in period 3. Which element is it?
Answer and reasoning
ANa A student who thinks electrons in the same shell need equal energies, so that any rise shows a new shell, picks this, because the second value is about twice the first. Rises of this size occur between valence electrons; the decisive jump, more than five times the value before it, comes at the third ionization.
BAl A student who thinks the jump comes when the last valence electron is removed picks this, counting three valence electrons because the jump is at the third ionization. The third electron is the first core electron, so only two electrons are valence electrons.
CMgCorrect The first two values rise modestly (738 to 1,451 kJ/mol), then the third is more than five times the second (7,733 kJ/mol). The jump marks the first core electron, so the atom has two valence electrons: a period 3 element in group 2, magnesium.
DSi A student who takes the largest value in the table as the first core electron picks this, counting four electrons before the fifth ionization. The values keep rising after the core is reached; the first core electron is found from the largest jump, at the third ionization.
Working Ratios of successive values: 1,451/738 = 2.0; 7,733/1,451 = 5.3; 10,543/7,733 = 1.4; 13,630/10,543 = 1.3. The large jump is between the 2nd and 3rd ionizations, so 2 valence electrons; period 3, group 2 = Mg.
A student wants to find out which of two elements, X or Y, has the greater first ionization energy, using photoelectron spectra of gaseous atoms of each element. Which comparison should the student make?
Answer and reasoning
ACompare the binding energies of the peaks of highest energy in the two spectra A student who reads the left-hand peak as the first electron removed picks this. The peak at the highest binding energy belongs to the 1s electrons, the most tightly held, which are not the electrons removed in the first ionization.
BCompare the heights of the peaks at the lowest binding energy in the two spectra A student who thinks a taller peak shows electrons that need more energy picks this. The height of a peak shows the number of electrons in that subshell; the energy is given by its position.
CCompare the combined heights of all of the peaks in each of the two spectra A student who thinks an atom with more protons always holds its outer electrons more strongly picks this, using the total number of electrons to compare nuclear charge. Atoms with more electrons often have lower first ionization energies, as potassium does compared with neon.
DCompare the binding energies of the lowest-energy peaks in the two spectraCorrect The first ionization energy is the energy needed to remove the most loosely held electron, which gives the peak at the lowest binding energy. The element whose lowest-energy peak lies at the higher binding energy has the greater first ionization energy.
A student draws the model shown to represent the relative sizes of sodium, magnesium and aluminum atoms. Which statement correctly evaluates the model?
Answer and reasoning
AIt is inconsistent: the nuclear charge rises while the valence electrons stay in the third shell, so the atoms should get smaller.Correct From Na to Al the number of protons increases from 11 to 13, while the valence electrons all occupy the third shell and shield one another poorly. The effective nuclear charge acting on them increases and pulls them closer, so the atoms get smaller: the model shows the opposite trend.
BIt is consistent: each atom has more electrons than the one before, and more electrons take up more space. A student who thinks more electrons make an atom larger picks this. The added electrons go into the same shell while the added protons pull all of the electrons closer, so the atoms get smaller across the period.
CIt is inconsistent: the three atoms should be the same size, because each added electron cancels an added proton. A student who thinks each added electron cancels the effect of an added proton picks this. Electrons in the same shell shield one another only weakly, so the effective nuclear charge increases from Na to Al and the atoms get smaller, not equal.
DIt is consistent: each atom is heavier than the one before, and heavier atoms take up more space than lighter ones. A student who thinks heavier atoms are larger picks this. Size depends on the attraction for the outer electrons and the shell they occupy, not on mass; Al atoms are heavier but smaller than Na atoms.
Which statement correctly compares the radius of a chloride ion, Cl⁻, with the radius of a chlorine atom, Cl, and explains the difference?
Answer and reasoning
ACl⁻ is smaller, because the extra electron gives the nucleus more negative charge to pull inward. A student who thinks an anion is pulled in because the nucleus has more charge to attract picks this. The nuclear charge is unchanged; the extra electron adds repulsion, so the anion is larger.
BCl⁻ is larger, because the added electron increases electron repulsion while the nuclear charge is unchanged.Correct Cl⁻ has 18 electrons and Cl has 17, both attracted by 17 protons. The extra electron in the third shell adds electron-electron repulsion without any extra nuclear charge, so the electron cloud spreads out and the ion is larger than the atom.
CThey are the same size, because both have three occupied shells around a nucleus of 17 protons. A student who thinks size is set only by the nucleus and the number of shells picks this. The number of shells is the same, but the extra electron increases repulsion in the outer shell, so Cl⁻ is larger than Cl.
DCl⁻ is smaller, because its complete octet is a stable arrangement that holds all of the electrons in more tightly. A student who thinks a filled octet holds electrons tightly because it is stable picks this. Completing the octet does not increase the nuclear charge; the added electron increases repulsion, so Cl⁻ is larger than Cl.
Which statement correctly compares the energy released when a gaseous S atom gains an electron with the energy released when a gaseous Cl atom gains an electron, and explains the difference?
Answer and reasoning
AMore is released for S, because Cl has more electrons, so its atom is larger and the added electron is farther away. A student who thinks an atom with more electrons is larger picks this. Cl atoms are smaller than S atoms, because the extra proton pulls the electrons of the same shell closer, so the electron added to Cl is attracted more strongly.
BMore is released for Cl, because gaining one electron gives a Cl atom a complete and stable octet. A student who thinks atoms attract electrons because of the stability of a full octet picks this. The comparison is right, but the energy released depends on the Coulombic attraction for the added electron, which is greater in Cl because of its greater nuclear charge acting at a shorter distance.
CThe same amount is released for both, because the extra proton in Cl is cancelled out by the extra electron it has. A student who thinks each added electron cancels an added proton picks this. Electrons in the same shell shield one another poorly, so the effective nuclear charge acting on the added electron is greater for Cl.
DMore is released for Cl, because the added electron enters the same shell but is attracted by one more proton.Correct In both atoms the added electron goes into the third shell, shielded by the same 10 inner electrons. The Cl nucleus has one more proton, so the effective nuclear charge acting on the added electron is greater and more energy is released (about 349 kJ/mol for Cl, compared with about 200 kJ/mol for S).
Which element, oxygen or sulfur, has the greater electronegativity, and what is the best justification?
Answer and reasoning
AOxygen, because the electrons it shares are closer to its nucleus and less shielded.Correct O and S are both in group 16. Oxygen's valence shell is the second shell, closer to the nucleus and shielded by only two inner electrons, so shared electrons are attracted more strongly by an oxygen atom (electronegativity 3.44 for O, 2.58 for S).
BSulfur, because its nucleus has more protons to attract the electrons that it shares in a bond. A student who thinks more protons always means a stronger pull picks this. Sulfur's shared electrons are in the third shell, farther from the nucleus and shielded by ten inner electrons, so they are attracted less strongly than oxygen's.
CSulfur, because its larger atom has more room in its outer shell for the electrons that it shares. A student who thinks a larger atom attracts electrons more strongly because it has more room picks this. Electrons in a larger atom's outer shell are farther from the nucleus and so are attracted less strongly.
DNeither, because all of the elements in the same group have equal electronegativities. A student who thinks elements in a group have the same values picks this. Electronegativity decreases down a group as the valence shell gets farther from the nucleus.
Which of the following elements has the greatest electronegativity?
Answer and reasoning
ANeon, Ne A student who thinks noble gases are the most electronegative because their shells are full picks this. Electronegativity describes attraction for shared electrons in a bond; neon forms no bonds and is not assigned a value, and fluorine is the most electronegative element.
BFluorine, FCorrect Fluorine has the greatest electronegativity of all the elements (3.98). Among elements that form bonds, the effective nuclear charge acting on the valence electrons is greatest at the right of a period, and the distance is smallest at the top of a group.
CSodium, Na A student who thinks a larger atom attracts shared electrons more strongly because it has more room in its outer shell picks this, the largest atom listed. Sodium's valence electron is in the third shell, shielded by ten inner electrons and attracted by an effective nuclear charge of only about 1+, so sodium has a low electronegativity (0.93).
DChlorine, Cl A student who thinks more protons always means a stronger pull picks this. Chlorine's valence electrons are in the third shell, farther from the nucleus than fluorine's, so chlorine (3.16) is less electronegative than fluorine.
The table gives the period, group and first ionization energy of magnesium and potassium, and the period and group of calcium. Which value is the best estimate of the first ionization energy of calcium?
Answer and reasoning
AAbout 420 kJ/mol A student who thinks elements in the same period have similar properties picks this, matching calcium to potassium, its neighbor in period 4. Calcium has one more proton than potassium acting on electrons in the same shell, so its first ionization energy is higher.
BAbout 590 kJ/molCorrect Calcium is below magnesium in group 2: its valence electrons are in the fourth shell, farther from the nucleus and more shielded, so its value is lower than 738 kJ/mol. It is to the right of potassium in period 4, with one more proton acting on electrons in the same shell, so its value is higher than 419 kJ/mol. Only 590 kJ/mol lies between the two (the measured value is 590 kJ/mol).
CAbout 740 kJ/mol A student who thinks elements in the same group have the same values picks this, giving calcium magnesium's value. Down a group the valence electrons are farther from the nucleus, so the first ionization energy decreases.
DAbout 900 kJ/mol A student who thinks more protons always mean a stronger pull on the outer electrons picks this, placing calcium above magnesium. Calcium's valence electrons are in a higher shell, so it holds them less strongly than magnesium does.
Working Ca is below Mg in group 2 (outer electrons in shell 4, farther and more shielded): IE(Ca) < 738 kJ/mol. Ca follows K in period 4 (one more proton, same shell): IE(Ca) > 419 kJ/mol. Of the options only about 590 kJ/mol satisfies 419 < IE < 738 (measured: 590 kJ/mol).
Compiled from the AP Chemistry Course and Exam Description (effective Fall 2024) and our question bank · Specialist review in progress. How these pages are made · Free, no account
