MCAT: General Chemistry
The valence electron in a lithium atom jumps from energy level n = 2 to n = 4. What is the energy of this transition in joules? In eV? (Note: RH = 2.18 x 10^18 J/electron = 13.6 eV/electron)
- 4.09 x 10^-19 J = -2.55 eV; energy is absorbed (calculations on pg 30)
In nuclear medicine, isotopes are created and used for various purposes; for instance, 18O is created from 18F. Determine the number of protons, neutrons and electrons in each of these species.
18O: 8 protons, 10 neutrons, 8 electrons 18F: 9 protons, 9 neutrons, 9 electrons
Calculate and compare the subatomic particles that make up the following atoms 19O 16O 17O 19F 16F 238U 240U
19O: 8 protons, 11 neutrons, 8 electrons 16O: 8 protons, 8 neutrons, 8 electrons 17O: 8 protons, 9 neutrons, 8 electrons 19F: 9 protons, 10 neutrons, 9 electrons 16F: 9 protons, 7 neutrons, 9 electrons 238U: 92 protons, 146 neutrons, 92 electrons 240U: 92 protons, 148 neutrons, 92 electrons
If given the following quantum numbers, which element(s) do they likely refer to? (Note: Assume that these quantum numbers describe the valence electrons in the element.) n = 2, l = 1 n = 3, l = 0 n = 5, l = 3 n = 4, l = 2
2p: B, C, N, O, F, Ne 3s: Na, Mg 5f: Actinide series 4d: Y, Zr, Nb, Mo, Tc, Ru, Rh, Pd, Ag, Cd
Calculate the energy of a photon of wavelength 662 nm. (Note: h = 6.626 x 10^-34 J*s)
3.00 x 10^-19 J (calculations on pg 30)
If an electron emits 3 eV of energy, what is the corresponding wavelength of the emitted photon? (Note: 1 eV = 1.60 x 10^19 J, h = 6.626 x 10^-34 J*s)
4.14 x 10^-7 m = 414 nm (calculations on pg 30)
Determine the number of protons, neutrons and electrons in a nickel-58 and in a nickel-60 +2 cation.
58Ni has an atomic number of 28 and a mass number of 58. Therefore, 58 Ni will have 28 protons, 28 electrons and 58-28, or 30, neutrons. 60Ni2+ has the same number of protons as the neutral 58 Ni atom. however, 60Ni2+ has apositive charge because it has lost two electrons; thus, Ni2+ will have 26 electrons. Also, the mass number is two units higher than for the 58Ni atom, and this difference in mass must be due to two extra neutrons; thus, it has a total of 32 neutrons.
Diamagnetism
A condition that arises when a substance has no unpaired electrons and is slightly repelled by a magnetic field.
Paramagnetism
A condition that arises when a substance has unpaired electrons and is slightly attracted to a magnetic field.
Period
A horizontal row of the periodic table containing elements with the same number of electron shells.
Conductor
A material in which electrons are able to transfer energy in the form of heat or electricity.
Electronegativity
A measure of the ability of an atom to attract the electrons in a bond; commonly measured with the Pauling scale.
Malleability
A physical property of metals that defines how well an element can be shaped using a hammer.
Lyman series
A portion of the emission spectrum for hydrogen representing electronic transitions from energy levels n > 1 to n = 1
Balmer series
A portion of the emission spectrum for hydrogen representing electronic transitions from energy levels n > 2 to n = 2.
Orbitals
A region of electron density around an atom or molecule containing no more than two electrons of opposite spin.
Octet rule
A rule stating that bonded atoms tend to undergo reactions that will produce a complete octet of valence electrons; applies without exception only to C, N, O and F.
Nucleus
A small central region of an atom; a dense, positively charged area containing protons and neutrons.
Neutrons
A subatomic particle contained within the nucleus of an atom; carries no charge and has a mass very slightly larger than that of a proton.
Protons
A subatomic particle that carries a single positive charge and has a mass slightly less than 1 amu.
Electrons
A subatomic particle that remains outside the nucleus and carries a single negative charge; in most cases, its mass is considered to be negligible.
Atomic mass unit (amu)
A unit of mass defined as 1/12 the mass of a carbon-12 atom; approx equal to the mass of one proton or one neutron.
Group
A vertical column of the periodic table containing elements that are similar in their chemical properties; also called a family.
What is the electron configuration for Zn+2?
A. 1s2 2s2 3s2 3p6 4s0 3d10 Remember that when electrons are removed from an element, forming a cation, they will be removed from the subshell with the highest n value first. Zn0 has 30 electrons, so it would have an electron configuration of 1s2 2s2 3s2 3p6 4s2 3d10. The 4s subshell has the highest principal quantum number, so it is emptied first, forming 1s2 2s2 3s2 3p6 4s0 3d10.
How many total electrons are in a 133Cs cation?
A. 54 The quickest way to solve this problem is to use the periodic table and find out how many protons are in Cs atoms; there are 55. Neutral Cs atoms would also have 55 electrons. A stable Cs cation will have a single positive charge because it has one unpaired s-electron. This translates to one fewer electron than the number of protons, or 54 electrons.
Which of the isotopes of carbon is LEAST likely to be found in nature? A. 6C B. 12C C. 13C D. 14C
A. 6C Recall that the superscript refers to the mass number of an atom, which is equal to the number of protons plus the number of neutrons present in an element. Sometimes a text will list the atomic number, Z, as a subscript under the mass number, A. According to the periodic table, carbon contains six protons, but differ in the number of neutrons. almost all atoms with Z> 1 have at least one neutron. Carbon is most likely to have a mass number of 12, for six protons and six neutrons.
Suppose that an atom fills its orbitals as stated: 3s: paired, 3p: 3 electrons unpaired Such an electron configuration most clearly illustrates which of the following laws of atomic physics? A. Hund's rule B. Heisenberg uncertainty principle C. Bohr model D. Rutherford model
A. Hund's rule The MCAT covers the topics in this chapter qualitatively more often than quantitatively. Its is critical to be able to distinguish the fundamental principles that determine electron organization, which are usually known by the names of the scientists who discovered or postulated them. The Heisenberg uncertainty principle refers to the inability to know the momentum and position of a single electron simultaneously. The Bohr model was an early attempt to describe the behavior of the single electron in a hydrogen atom. The Rutherford model describes a dense, positively charged nucleus. The element shown here, phosphorus, is often used to demonstrate Hund's rule because it contains a half-filled p subshell. Hund's rule explains that electrons fill empty orbitals first before doubling up electrons in the same orbital.
Consider the two sets of quantum numbers stated, which describe two different electrons in the same atom. n = 2, l = 1, ml = 1, ms = +1/2 n = 3, l = 1, ml = -1, ms = +1/2 Which of the following terms best describes these two electrons? A. Parallel B. Opposite C. Antiparallel D. Paired
A. Parallel The terms in the answer choices refer to the magnetic spin of the two electrons. The quantum number ms represents this property as a measure of an electron;s intrinsic spin. These electrons' spins are parallel, in that their spins are aligned in the same direction (ms = +1/2 for both species.
Which of the following are representative elements (A), and which are nonrepresentative (B)? Ag: Pb: Cu: N:
Ag: B Pb: A Cu: B N: A
Classify the following elements as metals (M), nonmetals (NM), or metalloids (MO): Ag: Pb: Cu: Si:
Ag: M Pb: M Cu: M Si: MO
Valence electrons
An electron in the highest occupied energy level of an atom; the tendency of a given valence electron to be retained or lost determines the chemical properties of an element.
Excited state
An electronic state having a higher energy than the ground state; typically attained by the absorption of a photon of a certain energy.
Metalloid
An element possessing properties intermediate between those of a metal and those of a nonmetal; also called a semimetal.
Anion
An ionic species with a negative charge.
Cation
An ionic species with a positive charge.
Transition metal
Any of the elements in the B groups of the periodic table, all of which have partially filled d subshells.
What are the definitions of atomic mass and atomic weight?
Atomic mass: Just slightly less than the sum of the masses of protons and neutrons in a given atom of an element. AToms of the same element with different mass numbers are isotopes of each other. Atomic weight: The weighted average of the naturally occurring isotopes of an element.
Mendeleev's table was arranged by atomic weight, but the modern periodic table is arranged by:
Atomic number
Isotopes
Atoms containing the same number of protons but different numbers of neutrons.
What equation describes the maximum number of electrons that can fill a subshell?
B. 4l + 2 This formula describes the number of electrons in terms of the azimuthal quantum number l, which ranges from 0 to n-1, with n being the principal quantum number.
The atomic weight of hydrogen is 1.008 amu. What is the percent composition of hydrogen by isotope, assuming that hydrogen 's only isotopes are 1H and 2D?
B. 99.2% H, 0.8% D The easiest way to approach thsi problem is to set up a system of two algebraic equations, where H and d are the percentages of H (mass = 1 amu) and d (mass = 2 amu), respectively. Your setup should look like the following system: H + D = 1 (percent H + percent D) = 100%) 1H + 2D = 1.008 (atomic weight calculation) Rearranging the first equation and substituting into the second yields (1-D) + 2D = 1.008, or D = 0.008. 0.008 is 0.8%, so there is 0.8% D.
Suppose an electron falls from n = 4 to its ground state, n = 1. Which of the following effects is most likely? A. A photon is absorbed. B. A photon is emitted. C. The electron moves into a p-orbital. D. The electron moves into a d-orbital.
B. A photon is emitted. Because the electron is moving into the n = 1 shell, the only subshell available is the 1s subshell. There will be some energy change, however, as the electron must lose energy to return to the minimum energy ground state. That will require emitting radiation in the form of a photon.
Which of the following transitions would result in the greatest gain in energy from a single hydrogen electron? A. An electron moving from n = 6 to n = 2. B. An electron moving from n = 2 to n = 6. C. An electron moving from n = 3 to n = 4. D. An electron moving from n = 4 to n = 3.
B. An electron moving from n = 2 to n = 6. For the electron to gain energy, it must absorb energy from photons to jump up to a higher energy level. There is a bigger jump between n = 2 and n = 6 than there is between n = 3 and n = 4.
Which of the following species is represented by the electron configuration 1s2 2s2 2p6 3s2 3p6 4s1 3d5? I. Cr II. Mn+ III. Fe2+
B. I and II when dealing with ions, you cannot directly approach electronic configuration based in the number of electrons they currently hold. First examine the neutral atom's configuration, and then demonstrate which electrons are removed. Due to the stablity of half-filled d-orbitals, neutral chromium assumes the electrons configuration of [AR]4s1 3d5. Mn must losed one electrons from its initial configuration to become the Mn+ cation. That electron would come from the 4s subshell according to the rule that the first electron removed comes from the highest energy shell. Fe must lose two electrons to become Fe2+. They'll both be lost from the same orbital; the only way Fe2+ could hold the configuration in the question stem would be if one d-electron and one s-electron were lost together.
Lithium and sodium have similar chemical properties. For example, both can form ionic bonds with chloride. Which of the following best explains this similarity? A. Both lithium and sodium ions are positively charged. B. Lithium and sodium are in the same group of the periodic table. C. Lithium and sodium are in the same period of the periodic table. D. Both lithium and sodium have low atomic weights.
B. Lithium and sodium are in the same group of the periodic table. The periodic table is organized into periods (rows) and groups (columns). Groups (columns) are particularly significant because they represent sets of elements with the same valence electron configuration, which in turn will dictate many of the chemical properties of these elements. The fact that both elements are positively charged does not explain the similarity in chemical properties; most metals produce positively charged ions.
Which of the following quantum sets is possible? A. n = 2; l = 2; ml = 1; ms = +1/2 B. n = 2; l = 1; ml = -1; ms = +1/2 C. n = 2; l = 0; ml = -1; ms = -1/2 D. n = 2; l = 0; ml = 1; ms = -1/2
B. n = 2; l = 1; ml = -1; ms = +1/2 The azimuthal quantum number l cannot be higher than n-1, ruling choice (A). The l number, which describes the chemical's magnetic properties, can only be an integer value between -l and l. It cannot be equal to +/- 1 id l= 0; this would imply that an s orbital has three subshells (-1, 0, and 1) when we know it only have one.
Rank the following elements by increasing electron affinity: Barium (Ba), copper (Cu), sulfur (S), and yttrium (Y).
Barium < yttrium < copper < sulfur
Magnetic resonance angiography (MRA) is a technique that can resolve defects like stenoic (narrowed) arteries. A contrast agent likely gadolinium or manganese injected into the blood stream interacts with the strong magnetic fields of the MRI device to produce such images. Based on their orbital configuration, are these contrast agents paramagnetic or diamagnetic?
Both of these molecules have unfilled valence electron shells with relatively few paired electrons; therefore, they are paramagnetic.
What is the maximum number of electrons allowed in a single atomic energy level in terms of the principal quantum number n?
C. 2n^2 For any value of n there will be a maximum of 2n^2 electrons; that is, two per orbital. This can also be determined from the periodic table. There are only two elements (H and He) that have valence electrons in the n=1 shell. Eight elements (Li to Ne) have valence electrons in the n = 2 shell. This is the only equation that matches this pattern.
Which of the following nest explains the inability to measure position and momentum exactly and simultaneously according to the Heisenberg uncertainty principle? A. Imprecision in the definition of the meter and kilogram. B. Limits on accuracy of existing scientific instruments. C. Error in one variable is increased by attempts to measure the other. D. Discrepancies between the masses of nuclei and of their component particles.
C. Error in one variable is increased by attempts to measure the other. The limitations placed by the Heisenberg uncertainty principle are caused by limitations inherent in the measuring process: if a particle is moving, it has momentum, but trying to measure that momentum necessarily creates uncertainty in the position.
Rank the following elements by decreasing first ionization energy: Calcium (Ca), carbon (C), germanium (Ge) and potassium (K)
Carbon > geranium > calcium > potassium
Which subatomic particle is the most important for determining each of the following properties of an atom? Charge: Atomic number: Isotope:
Charge: Determines by the number of electrons present. Atomic number: Determined by the number of protons. Isotope: Determined by the number of neutrons (while protons make up part of the mass number, it is the number of neutrons that explains the variability between isotopes).
An electron returns from an excited state to its ground state, emitting a photon at wavelength = 500 nm. What would be the magnitude of the energy change of one mole of these photons were emitted? (Note: h = 6.626 x 10^-34 J*s)
D. 2.39 x 10^5 J while daunting at first, the problem requires the MCAT favorite equation E = hc/lambda, where h = 6.626 x 10^-34 J*s (Planck's constant), c = 3.00 x 10^8 m/s is the speed of light, and lambda is the wavelength of the light. This question asks for the energy of one mole of photons, so we must multiply by Avogadro's number, NA = 6.02 x 10^23 mol^-1.
Which of the following atoms only has paired electrons in its ground state? A. Sodium B. Iron C. Cobalt D. Helium
D. Helium The only answer choice without unpaired electrons in its ground state is helium. Recall from the chapter that a diamagnetic substance is identified by the lack of unpaired electrons in its shell. A substance without unpaired electrons, like helium, cannot be magnetized by an external magnetic field and is actually slightly repelled. Elements that come at the end of a block (Group IIA, the group containing Zn, and the noble gases, most notably) have only paired electrons.
Alkali metals
Elements found in Group IA of the periodic table; highly reactive, readily losing one valence electron to form ionic compounds with nonmetals.
Alkaline earth metal
Elements found in Group IIA of the periodic table; chemistry is similar to that of alkali metals, except that they have two valence electrons and, thus, for 2+ cations.
Chalcogens
Elements found in group VIA of the periodic table with diverse chemistry; the group contains metals, nonmetals (like oxygen), and metalloids; typically form -2 anions.
Representative element
Elements in group 1, 2, and 13 through 18 in the modern IUPAC table ( the s- and p-blocks of the table, also called A group elements); these elements tend to have valence shells that follow the octet rule.
Nonrepresentative element
Elements with an expanded valence shell that includes d- and f-block electrons; also called Group B or transition elements.
Which will fill first, the 5d subshell or the 6s subshell?
For 5d, n = 5 and l = 2, so n + l = 7. For 6s, n = 6 and l = 0, so n + l = 6. Therefore, the 6s subshell has lower energy and will fill first.
For each of the properties below, write down the groups of the periodic table that exhibit those properties. High reactivity to water: Six valence electrons: Contain at least one metal: Multiple oxidation states: Negative oxidation states: Possess a full octet in the neutral state:
High reactivity to water: Groups 1 and 2 Six valence electrons: Groups 6 and 16 Contain at least one metal: Groups 1-15 Multiple oxidation states: All groups; most notably Group 3 through 12 (transition metals) Negative oxidation states: Almost all groups; most notably Group 14 through 17 (nonmetals) Possess a full octet in the neutral state: Group 18
Identify what type of element likely possesses the following properties: Luster: Poor conductivity of heat and electricity: Good conductivity but brittle:
Luster: Metals Poor conductivity of heat and electricity: Nonmetals Good conductivity but brittle: Metalloids
Write out and compare an orbital diagram for a neutral oxygen (O) atom and an O2- ion.
Neutral: 2s2 2p4 Ion: 2s2 2p6 Both O and O2- has fully fu=illed 1s- and 2s-orbitals. O has four electrons in the 2p subshell; two are paired, and the other two each have their own orbitals. O2- has six electrons in the 2p subshell, all of which are paired in the three p-orbitals.
According to Hund's rule, what are the orbital diagrams for nitrogen and iron?
Nitrogen has an atomic number of 7. Thus, its electron configuration is 1s2 2s2 2p3. According to Hund's rule, the two s-orbitals will fill completely, while the three p-orbitals will each contain one electron, all with parallel spins. Iron has an atomic number of 26. As determined earlier, its electron configuration is [Ar] 4s2 3d6. The electrons will fill all of the subshells except for 3d, which will contain four orbitals with parallel (upward) spin and one orbital with electrons of both spin directions. Subshells may be listed either in order which they fill (4s before 3d) or with subshells of the same principal quantum number grouped together, as shown here. Both methods are correct.
Metal
One of a class of elements on the left side of the periodic table possessing low ionization energies and electronegativities; readily give up electrons to form cations and possess relatively high electrical conductivity.
Nonmetal
One of a class of elements with high ionization energies and electron affinities that generally gain electrons to form anions; located in the upper right corner of the periodic table.
Rank the following elements by decreasing electronegativity: antimony (Sb), neon (Ne), oxygen (O), thallium (TI)
Oxygen > antimony > thallium > neon
Determine how many valence electrons come from each subshell in the following atoms: P in PO4^3- O in PO4^3- Ir Cf
P in PO4^3-: 2s, 6p, 2d, 10 total O in PO4^3-: 2s, 6p, 8 total Ir: 2s, 7d, 9 total Cf: 2s, 10f, 12 total
Halogens
The active nonmetals in Group VIIA of the periodic table, which have high electronegativities and high electron affinities.
An element Q consists of three different isotopes: A, B and C. Isotope A has an atomic mass of 40 amu and accounts for 60% of naturally occurring Q. Isotope B has an atomic mass of 44 amu and accounts for 25% of Q. Finally, isotope C has an atomic mass of 41 amu and accounts for 15% of Q. What is the atomic weight of element Q?
The atomic weight is the weighted average of the naturally occurring isotopes of that element: 0.60(40 amu) + 0.25(22 amu) + 0.15(41 amu) = 24.00 amu + 11.00 amu + 6.15 amu = 41.15 amu
Atomic radius
The average distance between a nucleus and its outermost electron; usually measured as one-half the distance between two nuclei of an element in its elemental form.
Ionic radius
The average distance from the center of the nucleus to the edge of its electron cloud; cationic radii are generally smaller than their parent metal, whereas anionic radii are generally larger than their parent nonmetal.
Effective nuclear charge (Zeff)
The charge perceived by an electron from the nucleus; applies most often to valence electrons and influences periodic trends such as atomic radius and ionization energy.
Aufbau principle
The concept that electrons fill energy levels in order of increasing energy, completely filling one sublevel before beginning to fill the next.
Heisenburg uncertainty principle
The concept that states that it is impossible to determine both the momentum and position of an electron simultaneously with perfect accuracy.
Subshells
The division of electron shells or energy levels into different values of the azimuthal quantum numbers (s, p, d and f); composed of orbitals.
What is the electron configuration of Fe3+?
The electron configuration of iron is [Ar] 4s2 3d6. Electrons are removed from the 4s subshell before the 3d subshell because it has a higher principal quantum number. Therefore, Fe3+ has a configuration of [Ar] 3d5, not [Ar] 4s2 3d3
Inert gases
The elements in Group VIIIA, which contain a full octet of valence electrons in their outermost shells, and are therefore very unreactive; also called noble gases.
Ionization energy (IE)
The energy required to remove an electron from the valence shell of a gaseous atom.
Principal quantum number (n)
The first quantum number, which defines the energy level or shell occupied by an electron
Spin quantum number (ms)
The fourth quantum number, which indicates the orientation of the intrinsic spin of an electron in an atom; can only assume values of +1/2 and -1/2.
Periodic law
The law stating that the chemical properties of elements depend on the atomic number of the elements and change in periodic fashion.
Atomic mass
The mass of a given isotope of an element; closely related to the mass number.
What is the electron configuration of osmium (Z = 76)?
The noble gas that comes just before osmium is xenon (Z = 54). Therefore, the electron configuration can begin with [Xe]. Continuing across the periodic table, we pass through the 6s subshell (cesium and barium), the 4f subshell (the lanthanide series; remember its position on the periodic table!), and the 5d subshell. Osmium is the sixth element in the 5d subshell, so the configuration is [Xe] 6s2 4f14 5d6.
Avogadro's number
The number of atomes or molecules in one mole of a substance: 6.02 x 10^23 mol^-1.
Atomic number (Z)
The number of protons in a given element
Pauli exclusion principle
The principle stating that no two electrons within an atom may have an identical set of quantum numbers.
Ductility
The property of metals that allows a material to be drawn into thinly stretched wires.
Azimuthal (angular momentum) quantum number (l)
The quantum number denoting the sublevel or subshell in which an electron can be found; reveals the shape of the orbital.
Hund's rule
The rule that electrons will fill into separate orbitals with parallel spins before pairing within an orbital.
Absorption spectrum
The series of discrete lines at characteristic frequencies representing the energy required to excite an electron from the ground state.
Electron configuration
The symbolic representation used to describe the electron arrangement within the energy sublevels in a given atom.
Magnetic quantum number (ml)
The third quantum number, defining the particular orbital of a subshell in which an electron resides; conveys information about the orientation of the orbital in space.
Ground state
The unexcited state of an electron.
Periodic Table of the Elements
The visual display of all known chemical elements arranged in rows (periods) and columns (groups) according to their atomic number and electron structure.
Atomic weight
The weighted average mass of the atoms of an element, taking into account the relative abundance of all naturally occurring isotopes.
While molar mass is typically written in grams per mole (g/mol), is the ratio moles per gram (mol/g) also acceptable?
This ratio is an equivalent concept. It is therefore acceptable, as long as units can be cancelled in dimensional analysis.
Paired electrons
Two electrons in the same orbital with assigned spins of +1/2 and -1/2.
Which electrons are the valence electrons of elemental vanadium, elemental selenium, and the sulfur atom in a sulfate ion?
Vanadium has five valence electrons: two in its 4s subshell and three in its 3d subshell. Selenium has six valence electrons: two in its 4s subshell and four in its 4p subshell. Selenium's 3d electrons are not part of its valence shell. Sulfur in a sulfur ion has 12 valence electrons: its original six plus six more from the oxygens to which it is bonded. Sulfur's 3s and 3p subshells can contain only eight of these 12 electrons; the other four electrons have entered the sulfur atom's 3d subshell, which is normally empty in elemental sulfur.
Rank the following elements by increasing atomic radius: niobium (Nb), praseodymium (Pr), tantalum (Ta) and xenon (Xe).
Xenon < niobium < tantalum < praseodymium