Introduction to General Chemistry

  1. All mass consists of tiny particles called atoms. Each atom is composed of a nucleus surrounded by one or more _____. The _____ of a nucleus is on the order of 10-4 angstroms (A). One angstrom is 10-10 m. The nucleus consists of _____ and ______, collectively called _____.
    • electrons
    • radius
    • protons and neutrons
    • nucleons
  2. Protons and neutrons are approximately equal in size and mass. Protons have a positive charge, while neutrons are electrically neutral. Protons and neutrons are held together to form the nucleus by the _____ _____ ____. The stability of the nucleus can be measured by its _____ _____ (define)
    • strong nuclear force
    • binding energy: the energy that would be required to break the nucleus into individual protons and neutrons.
  3. Surrounding the nucleus at a distance of about 1 to 3 A are electrons. The mass of an electron is more than _____ times smaller than the mass of a nucleon. Since the nucleons are ____ compared to the distance between the nucleus and the outermost electrons, the atom itself is composed mostly of ____ ____. If an atom were the size of a modern football stadium, it would have a nucleus the size of a _____. Matter is composed of atoms, therefore matter is mostly _____ ____.
    • 1800
    • tiny
    • empty space
    • marble
    • empty space
  4. State which of the three atom constituents is designated and state its charge
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  5. Electrons and protons have _____ charges of _____ magnitude. Although the charge on an electron is written as 1- and the charge on a proton is written as 1 +, remember that this charge is in electron units 'e' called the _____ ______. A charge of 1 e is equal to _______ coulombs, the SI unit for _____. When is an atom is electrically neutral?
    • opposite
    • equal electronic charge
    • 1.6 x 10-19
    • charge
    • When it contains the same number of protons and electrons.
  6. Elements are the building blocks of ______ and cannot be decomposed into simpler substances by _____ means. An element may have any number of ______ or ______, but only one number of _____.
    • compounds
    • chemical
    • neutrons or electrons
    • protons
  7. All elements can be depicted as shown in Figure 1.2. Z is the _____ number, indicating the number of protons. The atomic number provides the identity of the element because each element ,has a unique number of _____: if we know the atomic number, we know the _____ of the element.
    Label:
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    • atomic number
    • protons
    • identity
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  8. The mass number, A, is the number of ____ plus _____, and varies depending on the number of _____. Protons and neutrons each have a mass of approximately ___ amu and the mass of an atom is concentrated in the _____. This means that the mass number of an element is approximately equal to its ____ ____ or _____ _____.
    • protons
    • neutrons
    • neutrons
    • 1 amu
    • nucleus
    • atomic weight or molar mass
  9. The atomic weight given on the periodic table can be assigned either of the commonly used units: _____ _____ _____ (abbreviated as amu or the less commonly used SI abbreviation, u) or ____/____.
    • atomic mass units
    • grams/mole
  10. Why is it that not every atom of a given element will have a mass equal to the mass number? What is the atomic weight is equal to?
    Simply because isotopes of an atom have different numbers of neutrons.

    The atomic weight is equal to the weighted average of the naturally occurring isotopes of that element
  11. The nucleus of a specific isotope (define) is called a ____. Isotopes have similar _______ properties. Hydrogen has three important isotopes:____, ______, and _____ 99.98% of naturally occurring hydrogen is _____.
    Isotopes are two or more atoms of the same element that contain different numbers of neutrons.

    • nuclide
    • chemical
    • 1H (protium), 2H (deuterium), and 3H (tritium)
    • protium
  12. The isotopes for carbon include ____, ____, and ____. Each of carbon's isotopes contains ___ protons with _____ ____ of neutrons. The ____ protons are what define carbon; if the number of _____ changed, it would no longer be carbon. 12C (carbon-12) contains ___ neutrons, 13C (carbon-13) contains ___ neutrons , and 14C (carbon-14) contains ___ neutrons.
    • 12C, 13C, and 14C
    • 6 protons
    • varying numbers
    • six protons
    • protons
    • 6 neutrons
    • 7 neutrons
    • 8 neutrons
  13. Ions (define) are not electrically _____. Positive ions have fewer electrons than protons and are called _____; negative ions have more electrons than protons and are called _____. Define salt
    Ion: When the number of electrons in an atom does not equal the number of protons, the atom carries a charge and is called an ion

    • neutral
    • cations
    • anions

    Salt: A neutral compound composed of a positive and a negative ion together.
  14. When a neutral atom loses an electron to become a cation, it gets ______. The atom still has the same number of ____, so there are now more _____ than ______. The result is that the positive charge of the nucleus exerts a greater attractive force on each _____ _____, pulling them closer to the ______.
    • smaller
    • protons
    • protons
    • electrons
    • valence electron
    • nucleus
  15. The loss of an electron also ______ the repulsive forces between the electrons, further contributing to the _____ in size. On the other hand, when a neutral atom gains an electron to become an _____ , it gets ______. The atom now has more electrons than protons, so the positive charge of the nucleus pulls _____ strongly on each individual valence electron. The addition of an electron also ______ the repulsive forces between the electrons, pushing them farther away from each other. The net effect is that the anion is _____ than the neutral atom.
    • reduces
    • decrease
    • anion
    • larger
    • less
    • increases
    • larger
  16. The periodic table lists the elements from left to right in the order of their _____ numbers. Each horizontal row is called a _____. The vertical columns are called _____ or _____. As will be described in more detail below, elements in the same family share some similar ______ and ______ properties.
    • atomic number
    • period
    • groups or families
    • chemical and physical properties

    *An element's relative location on the periodic table can be used to predict and recall both the characteristics of that element and how it will interact with other elements and molecules. Understanding how the periodic table is organized minimizes the need to memorize and turns the periodic table into a key reference for solving many questions on the MCAT®.
  17. The periodic table in Figure 1.4 divides the elements into three sections:
    • 1. nonmetals on the right (green)
    • 2. metals in the middle and on the left (yellow)
    • 3. metalloids along the diagonal separating the metals from the nonmetals (blue).
  18. *There are two methods commonly used to number the groups. The newer method is to number them 1 through 18 from left to right. An older method that is still occasionally used is to separate the groups into sections A and B, and then number them with R oman numerals as show n in Figure 1.4. The periodic table on the MCAT® does not have group numbers.
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  19. Metals (define) can be described as atoms in a sea of ______, which emphasizes their loose hold on their electrons and the fluid-like nature of their _____ _____. What gives them (metals) their metallic character?
    • Metals are large atoms that tend to lose electrons to form positive ions and positive oxidation states. 
    • electrons
    • valence electrons
    • The easy movement of electrons within metals is what gives them their metallic character.
  20. 4 metallic characteristics of metals
    • 1)ductility (easily stretched) 
    • 2)malleability (easily hammered into thin strips)
    • 3)thermal and electrical conductivity
    • 4)a characteristic luster.
  21. Metal atoms easily ____ ____ each other, allowing metals to be hammered in to thin sheets or drawn into wires. _____ move easily from one metal atom to the next, transferring _____ or _____ in the form of _____ or _____. All metals except ______ exist as solids at room temperature. Metals typically lose ______ to become cations, and form _____ bonds.
    • slide past
    • Electrons
    • energy or charge
    • heat or electricity
    • mercury
    • electrons
    • ionic bonds
  22. Nonmetals have diverse _____ and ______ behaviors. Molecular substances are made from nonmetals, since nonmetals form _____ bonds with one another. Generally speaking, nonmetals have _____ melting points than metals. They tend to form _____, which commonly react with ____ _____ to form ionic compounds. Metalloids have some ______ and some ___-_____ characteristics.
    • appearances and chemical behaviors
    • covalent bonds 
    • lower
    • anions
    • metal cations
    • metallic
    • non-metallic
  23. The section A groups (groups 1, 2, 13, 14, 15, 16, 17, 18) are known as the _______ elements or main-group elements and the section B groups (groups 3, 4, 5, 6, 7, 8, 9, 10, 11, and 12) are called the _____ _____.
    • representative elements
    • transition metals
  24. It is important to be able to predict what kind of ion an element will form, i.e. whether it will gain or lose ______. The representative elements make ions by forming the closest _____ _____ _____ configuration; this is why metals tend to form cations and nonmetals tend to form _____.
    • electrons
    • noble gas electron
    • anions
  25. When the transition metals form ions, they lose electrons first from their _____ ______ and then from their ______, as will be discussed later in this lecture.
    • highest s-subshell
    • d-subshell
  26. The periodic table below shows some of the common ions formed by the _____ _____. Notice in the figure below that Group ___ elements make 1+ ions. There are five 3+ ions: ___, ___, ___, ___, and ___. The other transition metal ions are ___.
    • transition metals
    • 11
    • Cr3+, Fe3+, Au3+, Al3+, and Bi3+
    • 2+
  27. Easy here! You don't have to memorize the charge on every cation made by transition metals. This is background knowledge. But you should be able to predict the charge based upon two things:
    • 1. Atoms lose electrons from the highest energy shell first. This means that in transition metals, electrons are lost from the s-subshell first, and then from the d-subshell.
    • 2 . Ions seek symmetry. Representative elements form noble gas electron configurations when they make ions. For example, group 1 atoms form 1+ cations and group 17 atoms form 1- anions. Transition metals try to "even-out" their d-orbitals so that each orbital has the same number of electrons. Whenever possible, an ion will have a half filled or completely filled orbital. Important reference points in the periodic table include the periods 1 and 2, which have half-filled and completely filled s orbitals, periods 7 (VIIB) and 12 (IIB), which have half-filled and completely filled d orbitals, and periods 15 and 18, which have half-filled and completely filled porbitals.
  28. Elements in the same group on the periodic table have _____ chemical properties (explain)
    • similar
    • because they have the same number of valence electrons, or electrons in the outermost shell. They tend to make the same number of bonds and exist as similarly charged ions.
  29. The MCAT® requires the ability to identify and make predictions about specific groups: Group 1 (_____ ____), Group 2 (____ _____ ____), Group 16 (the ____ group), Group 17 (_____), and Group 18 (_____ ____). *This section will consider the properties that are useful for making predictions about elements in these groups.
    • alkali metals
    • alkaline earth metals
    • oxygen
    • halogens
    • noble gases
  30. State 6 characteristics of the Group 1 or alkali metals as pure substances
    • 1)They are soft metallic solids
    • 2)low densities
    • 3)low melting points
    • 4)They easily form 1 + cations, such as Na+.
    • 5)They are highly reactive, reacting with most nonmetals to form ionic compounds. Alkali metals react with hydrogen to form hydrides such as NaH. Alkali metals react exothermically (and explosively) with water to produce the respective metal hydroxide and hydrogen gas.
    • 6)In nature, alkali metals exist only in compounds.
  31. 6 characteristics of hydrogen
    • 1)Hydrogen is unique, as its chemical and physical characteristics do not conform well to any single family.
    • 2)It is a nonmetal and therefore can form covalent bonds.
    • 3)It can also form ionic compounds with metal cations as the anion hydride.
    • 4)Under most conditions, hydrogen is a colorless, odorless, diatomic gas.
    • 5)Hydrogen plays a significant role in multiple chemical and physical processes, particularly acid-base chemistry and intermolecular forces.
    • 6)Hydrogen was placed in the first column of the periodic table because it has one electron in the s-orbital, but it does not share its characteristics with this group.
  32. 5 characteristics of the alkaline earth metals (group 2)
    • 1)Harder, more dense, and melt at higher temperatures than alkali metals.
    • 2)They form 2+ cations, such as Mg2+.
    • 3)They are less reactive than alkali metals because their highest energy electron completes the s orbital.
    • 4)Heavier alkaline earth metals are more reactive than lighter ones.
    • 5)Like the alkali metals, the alkaline earth metals exist only in compounds in nature.
  33. 3 characteristics of group 14 elements
    • 1)All the Group 14 elements can form four covalent bonds with nonmetals.
    • 2)All beyond the second period can form two additional bonds with Lewis bases using d orbitals.
    • 3)Only carbon forms strong π-bonds to make strong double and even triple bonds. This characteristic of carbon is often critical to the structure of biological molecules
  34. 7 Characteristics of Group 15 elelements
    • 1)Group 15 elements can form 3 covalent bonds.
    • 2)In addition, all beyond the second period can form two additional covalent bonds by using their d orbitals.
    • 3)These elements can further bond with a Lewis base to form a sixth covalent bond.
    • 4)Nitrogen can form a fourth covalent bond by donating its lone pair of electrons to form a bond.
    • 5)Nitrogen forms strong π-bonds to make double and triple bonds.
    • 6)Phosphorous can form only weak π bonds to make double bonds.
    • 7)The other Group 15 elements do not make π-bonds.
  35. 9 characteristics of the Group 16 elements
    • 1)Group 16 elements are called the chalcogens or oxygen group.
    • 2)Oxygen and sulfur are the important chalcogens for the MCAT®. Oxygen is the second most electronegative element.
    • 3)It is divalent and can form strong n-bonds to make double bonds.
    • 4)In nature, oxygen exists as O2 (dioxygen) and O3 (ozone).
    • 5)Oxygen typically reacts with metals to form metal oxides. Alkali metals form peroxides (Na2O2) and super oxides (KO2) with oxygen.
    • 6)The most common form of pure sulfur is the yellow solid S8 .
    • 7)Metal sulfides, such as Na2S, are the most common form of sulfur found in nature.
    • 8)Sulfur can form two, three, four, five, or even six bonds.
    • 9)It has the ability to π-bond, forming strong double bonds.
  36. 8 characteristics of the group 17 elements
    • 1)The radioactively stable Group 17 elements, called halogens, are fluorine, chlorine, bromine, and iodine.
    • 2)Halogens are highly reactive.
    • 3)Fluorine and chlorine are diatomic gases at room temperature; bromine, a diatomic liquid; and iodine, a diatomic solid.
    • 4)Halogens like to gain an electron to attain a noble gas configuration. However, halogens other than fluorine can take on oxidation states as high as + 7 when bonding to other highly electronegative atoms such as oxygen.
    • 5)When in compounds, fluorine always has an oxidation state of -1. This means that fluorine can make only one bond.
    • 6)The other halogens can make more than one bond, though this is rare.
    • 7)All of the halogens can combine with hydrogen to form gaseous hydrogen halides. The hydrogen halides are soluble in water, forming the hydrohalic acids.
    • 8)Halogens react with metals to form ionic halides, such as NaCI.
  37. 4 characteristics of the group 18 elements
    • 1)The Group 18 noble gases (also called the inert gases) are nonreactive.
    • 2)Unlike other elements, the noble gases are normally found in nature as isolated atoms.
    • 3)They are all gases at room temperature.
    • 4)Noble gases are very stable. Representative elements form ions with the electron configurations of noble gases in order to gain stability.
  38. The elements that tend to exist as diatomic molecules are _____, _____, _____, and _____. When these elements are discussed, it is safe to assume that they are in their ______ form unless otherwise stated. In other words, the statement "nitrogen is nonreactive" refers to ___ rather than ___.
    • hydrogen, oxygen, nitrogen, and the halogens
    • diatomic
    • N2
    • N
  39. This section will describe how the periodic table can be used to predict changes in the chemical properties of atoms across a period or down a group. These predictions can be summarized as four periodic trends: _____ _____, _____ _____, ________ and _____ _____. As a general rule, the atomic radius increases going "____" and "to the ____" on the periodic table, while the other properties increase "___" and "to the _____."
    • atomic radius, ionization energy, electronegativity, and electron affinity
    • down
    • left
    • up
    • right
  40. The atomic radius is the distance from the center of the _____ to the outermost _____. Atomic radius corresponds to the size of the atom. Moving across a period, the radius ______ (explain)
    • nucleus
    • electron
    • decreases because each subsequent element has an additional proton, which pulls more strongly on the surrounding electrons.
  41. However, moving down a group, new shells of ______ are added. These outer electrons are "______" from the attraction of the _____ in the inner nucleus. As a result, atomic radius ______ going down a group.
    • electrons
    • shielded
    • protons
    • increases
  42. To gain a deeper understanding of changes in atomic radius, and to understand the other periodic trends, consider Coulomb's law, F =kq1q2/r2. Coulomb's law describes the _______ force (define) between an ______ and the _____.
    • electrostatic force: the force between charged objects, which is attractive between opposite charges and repulsive between like charges.
    • electron and the nucleus
  43. With respect to coulomb's law, the distance between the electron and the nucleus is __. For q1 we might plug in the ______ charge of the nucleus, (__), and for q2 , the charge on an ______, (__). This would work well for hydrogen, where the lone electron feels ___% of the positive charge on the nucleus.
    • r
    • positive
    • Z
    • electron (e)
    • 100%
  44. However, in helium the _____ electron shields some of the nuclear charge from the _____ electron, so that it doesn't feel the entire nuclear charge Z. Shielding occurs due to the repulsive forces between ______.
    • first
    • second
    • electrons
  45. The amount of charge felt by the most recently added electron is called the ______ _______ ______ (___)· In complete shielding, each ______ added to an atom would be completely shielded from the attractive force of all the _____ except for the _____ _____ added, and the Zeff would be ___ eV for each electron.
    • effective nuclear charge (Zeff)
    • electron
    • protons
    • last proton
    • 1 eV
  46. Without shielding, each electron added would feel the full ______ force of all the ______ in the nucleus, and the Zeff would simply be equal to ___ for each electron.
    • attractive
    • protons
    • Z
  47. Figure 1.6 shows Zeff values (given in electron-volts) for the highest energy electron in each element through sodium. Zeff generally _____ going left to right across the periodic table. While more protons are added across a period, ______ Z, the new electrons added are in Electron Salty roughly the same energy level, and therefore do not experience significantly more ______ than the previous electron.
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    • increases
    • increasing
    • shielding
  48. Zeff also _____ going from top to bottom down the periodic table. Though the energy level of the outermost electrons _____ down a group, the attractive pull of the growing positively charged nucleus ______ the additional shielding effects of higher electron shells.
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    • increases
    • increases
    • outweighs
  49. In Figure 1.6, notice that Zeff drops going from neon to sodium (explain). This causes a strong ______ in shielding and ______ in Zeff, but the outermost electron in sodium still experiences a _____ Zeff than the outermost electron of the element immediately above it on the periodic table, lithium (explain). A similar drop also occurred between He and Li, though it was not quite as large because there were fewer _____ and _____ involved.
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    This happens because the new electron is added to an entirely new shell, the 3s subshell.

    • increase
    • reduction
    • higher

    This is because the effects of the more strongly charged nucleus outweigh the shielding effect that an additional electron shell can provide.

    protons and electrons
  50. The ____, and not __, should be plugged in for q1 in Coulomb's law to find the force on the outermost electron. The force on an electron is a function of both q1 (___) and r (the distance from the ____). Notice that Zeff can be used to explain the atomic radius trend discussed earlier. (Demonstrate: 3-Story)
    • Zeff
    • Z
    • Zeff
    • nucleus

    • Since the effective nuclear charge increases from left to right on the periodic table, each additional electron is pulled more strongly toward the nucleus.
    • The result is that atoms tend to get smaller when adding electrons across the periodic table.
    • When moving down a group, each drop represents the addition of a new electron shell, and thus atoms tend to increase in size moving down a group.
  51. Zeff can also be used to understand trends in ion size. As discussed previously, cations are _____ than the neutral element while anions are _____. When an atom loses an electron, Zeff _____ because there are now more _____ relative to _____. ______Zeff means electrons are pulled closer to the nucleus. When an atom gains an electron, Zeff ______ because there are now more ______ relative to _____. Shielding increases due to the increased ______ _____ between electrons, and ______ are pushed further away from the nucleus.
    • smaller
    • larger
    • increases
    • protons
    • electrons
    • Increased
    • decreases
    • electrons
    • protons
    • repulsive forces
    • electrons
  52. Isoelectronic ions are ions with the same number of ______, but different ______ identities. O2-, F, neutral Ne, Na+, and Mg2+ all have the same number of electrons. However, since they have different numbers of protons, the electrons feel different ______ ______ _____. As a result, isoelectronic ions are not the same ____. The largest of these is ___. Its electrons feel the ______ Zeff (why?). The smallest of these is Mg2+. Its electrons feel the ______ Zeff (explain)
    • electrons
    • elemental
    • effective nuclear charges
    • size
    • O2-
    • weakest

    because there are two more electrons than protons.

    strongest

    because there are two more protons than electrons.
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  54. Ionization energy (define) generally increases along the periodic table from ____ to _____ and from _____ to ____. When an electron is more strongly attracted to the nucleus, _____ energy is required to detach it.
    Ionization energy: the energy needed to detach an electron from an atom.

    • left to right
    • bottom to top
    • more
  55. Define first and second ionization energy
    • The first ionization energy: the energy necessary to remove an electron from neutral atom in its gaseous state to form a +1 cation.
    • The second ionization energy: the energy required for the removal of a second electron from the same atom to form a +2 cation.
    • **Third, fourth, fifth, and other ionization energies are named in the same manner.
  56. The second ionization energy is always _____ than the first (explain)
    greater

    because once one electron is removed, the effective nuclear charge increases for the remaining electrons.
  57. For periodic trends, all the "E's" are up and to the right: (list them)
    What happens to the other trend
    • ionization Energy
    • Electronegativity
    • Electron affinity

    The other trend -atomic size- follows the opposite pattern (down and to the left)
  58. The ionization energy trend can be remembered by considering its relation to
    zeff and r. Moving across a period to the right, ______ zeff values pull electrons ____ strongly toward the nucleus. Therefore, _____ energy is required to rip them off
    • increasing
    • more
    • more
  59. Moving down a group, Zeff _____, but the distance of the electron from the nucleus _____ as well. Coulomb's law demonstrates that the electrostatic force, F, ______ with the square of the distance from the nucleus, r. Due to the exponent, the increased distance plays a ____ important role than the increased Zeff and the attractive electrostatic force ______ down a group. Less force means _____ energy is required to remove the electron, so ionization energy ______ moving down a group.
    • increases
    • increases
    • decreases
    • more
    • decreases
    • less
    • decreases
  60. Define electronegativity
    When two atoms have different electronegativities, they share ______ unequally, causing ______. Relative electronegativity determines the direction of ______ within a bond and within an overall molecule.
    Electronegativity: the tendency of an atom to attract electrons shared in a covalent bond.


    • electrons
    • polarity
    • polarity
  61. Like ionization energy, electronegativity tends to increase across a period from ____ to _____ and ____ a group. The most commonly used measurement of electronegativity is the ______ scale, which ranges from a value of 0.79 for cesium to a value of 4.0 for fluorine. It is important to remember that _____ is the most electronegative element.
    • left to right and up a group 
    • Pauling scale
    • fluorine
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  62. The electronegativity of hydrogen falls between that of _____ and that of _____. When bonded with hydrogen, carbon and elements to the right of carbon will carry a partial ______ charge while hydrogen will carry a partial ______ charge. Think of CH4. Boron and the elements to the left of boron will carry a partial ______ charge when bonded to hydrogen, while the hydrogen will carry a partial ______ charge. Think of the hydrides. (H-) in NaH or LiAIH3
    • boron
    • carbon
    • negative
    • positive
    • positive
    • negative
  63. Since noble gases tend not to make bonds, electronegativity values are ______ for the noble gases. Electronegativity values provide a system for predicting which type of bond will form between two atoms. Atoms with large differences in electronegativity (__ or larger on the Pauling scale as a rule of thumb) will form _____ bonds. Metals and non-metals usually exhibit ____ electronegativity differences and form ____ bonds with each other.
    • undefined
    • 1.6
    • ionic
    • large
    • ionic
  64. Atoms with moderate differences in electronegativities (___ -___ on the Pauling scale) will generally form _____ _____ ____. Atoms with very minor electronegativity differences (___ or smaller on the Pauling scale) will form ______ _____ bonds.
    • 0.5 - 1.5
    • polar covalent bonds
    • 0.4
    • nonpolar covalent
  65. Define Electron affinity
    Electron affinity: the willingness of an atom to accept an additional electron. More precisely, it is the energy released when an electron is added to an isolated atom.
  66. Just like electronegativity, electron affinity tends to increase on the periodic table from _____ to _____ and from ______ to ____. The sign of electron affinity values can be different for different atoms (explain)
    • left to right
    • bottom to top

    because some atoms release energy when accepting an electron (and thus become more stable), while others require energy input to force the addition of an electron (since the additional electron decreases stability).
  67. Warning: Electron affinity is sometimes described in terms of exothermicity, for which the energy released is given a negative sign. We can state this as follows: electron affinity is more exothermic to the right and up on the periodic table. The noble gases ____ follow this trend. Electron affinity values for the noble gases are ______ (why)
    • don't 
    • endothermic

    because noble gases are stable and thus significant amounts of energy are required to force them to take on electrons and become less stable.
  68. Neils Bohr proposed a theory of the atom, known as the Bohr atom, which represents the atom as a nucleus surrounded by electrons in discrete _____ _____. In the orbital structure of the hydrogen atom, a single electron orbits the hydrogen nucleus in an _____ _____. Any electron in any atom of any element, including this hydrogen electron, can be described by four ______ ______, which serve as that electron's "ID number." State the Pauli Exclusion Principle
    • electron shells
    • electron shell
    • quantum numbers
    • The Pauli Exclusion Principle states that no two electrons in the same atom can have the same four quantum numbers.
  69. The first quantum number is the ____ _____ _____ (__). (Define)
    principle quantum number (n): It designates the shell level of the electron, with low numbers closest to the nucleus.
  70. The second quantum number is the _______ ______ _____ (__). (Define)
    azimuthal quantum number (Q): It designates the electron's subs hell, each of which has a distinct shape. l = 0 is the s subshell, l = 1 is the p subshell, l= 2 is the d subshell, and l = 3 is the f subshell.
  71. The third quantum number is the _____ _____ ______ (__). (Define)
    magnetic quantum number (mQ):designates a precise orbital within a given subshell. Each orbital can hold two electrons. Each subshell has orbitals with magnetic quantum numbers ranging from -l to +l.
  72. Within the p-subshell (l = 1), there are three magnetic quantum numbers: ___, ___, and ___. An atom will spread out its electrons amongst the orbitals, such that all orbitals in a subshell have ___ electron before any has ___. The number of total orbitals within a shell is equal to ___. Solving for the number of orbitals for each shell gives 1, 4, 9, 16 ... Since there are ____ electrons in each orbital, the number of elements in the periods of the periodic table is ___, ___, ___, and ___.
    • -1, 0, and 1
    • one
    • two
    • n2
    • two
    • 2, 8, 18, and 32
  73. The fourth quantum number is the _____ _____ _____ _____ (__). The electron spin quantum number has possible values of___ or ___. This final quantum number is used to distinguish between the ___ electrons that may occupy the same orbital, since those electrons will have the same first ____ quantum numbers.
    • electron spin quantum number (m5)
    • -1/2 or +1/2
    • two
    • three
  74. The size of an atom has a significant effect on its chemistry. Small atoms hold charge in a concentrated way because they have _____ orbitals available to distribute and thereby ______ charge. This concentration of charge makes the ______ element in each group bonds _____ readily and with _____ bond strength, especially when in _____ form. Fluorine is a good example. The fluoride ion (F) is too _____ to manage its full negative charge. Therefore, it is generally insoluble and bonds immediately when in solution. For this reason, in toothpaste, fluoride (a poison in high concentrations) _____ immediately with the enamel of teeth before it can be ingested.
    • fewer
    • stabilize
    • smallest
    • more
    • greater
    • ionic
    • small
    • bonds
  75. Small atoms do not have d-orbitals available for bond formation, and therefore cannot form more than ___ bonds. Large atoms have d-orbitals, allowing for more than __ bonds. Oxygen typically forms two bonds, while the Larger sulfur can form up to ___. Smaller atoms have the advantage when it comes to __-bonding. The p-orbitals on large atoms do not _____ significantly, so they cannot easily form __-bonds. Carbon, nitrogen, and oxygen are small enough to form ____ π-bonds while their Larger third row family members form only ____ π-bonds, if they form π-bonds at all.
    • 4 bonds 
    • 4 bonds
    • six
    • π-bonding
    • overlap
    • π-bonds
    • strong
    • weak
  76. Let's summarize: The flrst quantum number is the ____. It corresponds roughly to the energy level of the _____ within that shell. The second quantum number is the ______. It gives the _____. Recognize that s orbitals are ______ and p orbitals are ______-shaped. The third quantum number'is the specific _____ within a subshell. The fourth quantum number distinguishes between two ______ in the same orbital; one is ____ ____ and the other is ____ ____.

    State the Name, symbol, value and character of the quantum numbers
    Image Upload 11
    • shell
    • electrons
    • subshell
    • shape
    • spherical
    • dumbbell-shaped
    • orbital
    • electrons
    • spin +1/2
    • spin -1/2

    Image Upload 12
  77. State Heisenberg's Uncertainty principle
    This uncertainty arises from the dual nature (wave-particle) of matter, and is on the order of Planck's constant (6.6 x 10-34 J s):
    The Heisenberg Uncertainty Principle states that there is an inherent uncertainty in the product of the position of a particle and its momentum.

    ΔxΔp≥h/2
  78. To put the uncertainty principle more simply, the more we know about the _______ of any particle, the less we can know about its ______, and vice versa. There are other quantities besides position and momentum to which the uncertainty principle applies, but position and momentum are the ones that will likely be tested on the MCAT®.
    • momentum
    • position
  79. Define the Aufbau principle
    The Aufbau principle: (sometimes called the "building up principle") states that with each new proton added to create a new element, the new electron that is added to maintain neutrality will occupy the lowest energy level available.
  80. All other things being equal, the lower the energy state of a _____, the more stable the system. Thus, electrons look for an orbital in the _____ energy level whenever they are added to an atom. The orbital with the ______ energy will be located in the ______ with the lowest energy.
    • system
    • lowest
    • lowest
    • subshell
  81. For the representative elements, the shell level of the most recently added electrons is given by the _____ in the periodic table. The most recently added electron for Sr is in shell __, and the most recently added electron for Sr is in shell __.
    • period
    • 2
    • 5
  82. For transition metals the shell of the most recently added electron lags ____ behind the period. The most recently added electron of Ag is in shell 4, though it is in the ____ period. For the lanthanides and actinides, the shell of the most recently added electron lags ____ behind the period. The most recently added electron for Ce is in shell 4, though the lanthanide series corresponds to the ____ period.
    • one
    • fifth
    • two
    • sixth
  83. The subshells (s, p, d, and}) are the orbital shapes with which we are familiar. Orbital shapes are not true paths that  electrons follow, but rather represent _____ _____ for the position of an _____. There is a 90% chance of finding the electron somewhere inside the given shape. Be familiar with the shapes of the orbitals in the __- and __-subshells. The shapes of the d- and f-subshells are beyond the scope of the MCAT®
    • probability functions
    • electron
    • s- and p- subshells
  84. The organization of the periodic table indicates where each type of _____ is filling. The s-subshells are filling in groups ___ and ___, the p-subshells are filling in groups ___-___, the d-subshells are filling in groups ___-___, and the f-subshells are filling in the _______ and ______ series. Therefore, the periodic table can be used to determine the _____ and _____ of the most recently added electron for any element. For example, the most recently added ______ in S is in the 3p energy level; for Au it is in the ___ energy level; and forK it is in the ___ energy level.
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    • subshell
    • 1 and 2
    • 13-18
    • 3-12
    • lanthanide and actinide series
    • shell and subshell
    • electron
    • 5d energy level 
    • 4s energy level
  85. Valence electrons, the electrons which contribute most to an element's ______ properties, are located in the ______ _____ of an atom. In most cases, only electrons from the ___ and ___ subshells are considered valence electrons.
    • chemical
    • outermost shell
    • s and p subshells
Author
chikeokjr
ID
339000
Card Set
Introduction to General Chemistry
Description
Chemistry Ch 1 (Pt I)
Updated