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Chem unit 4
Chem 7.5-9.3
| Question | Answer |
|---|---|
| What is an electron configuration? | shows how electrons are distributed across energy levels and sublevels (s, p, d, f) around the nucleus |
| What three quantum numbers are needed to describe an atomic orbital? | The principal quantum number (n), the angular moment quantum number (l), and the magnetic quantum number (ml) |
| What is the electron spin quantum number, Ms, and how does it arise? | describes the intrinsic spin of an e-; arises b/c e- have tiny mechanical magnets with an inherent angular momentum |
| What are the allowed values of Ms? | +1/2 and -1/2 |
| What is an orbital diagram? | visual way to show how electrons fill orbitals in an atom using boxes and arrows |
| What is the Pauli exclusion principle? | states that no two electrons in the same atoms can have the same set of four quantum numbers |
| What does it mean for electrons to be spin-paired? | electrons must share the same orbital and have opposite spins |
| In general, for a given shell of orbitals, orbitals with a larger l value has _________________energies | higher |
| How does Coulomb’s law contribute to relative orbital energies? | b/c orbitals with electron density closer to the nucleus are lower in energy b/c Coulomb's law makes the electron-nucleus attraction stronger |
| What is shielding and how does it contribute to relative orbital energies? | the reduction of the effective nuclear charge felt by an electron b/c other e- block some of the nucleus's positive charge |
| What is effective nuclear charge and why is it less than nuclear charge? | net positive charge that an e- actually feels from the nucleus after accounting for shielding by other electrons |
| What is penetration and how does it contribute to relative orbital energies? | describes how much an electron in a given orbital can get close to the nucleus |
| What is the general ordering of subshell energies in a given shell of orbitals? | s<p<d<f |
| Describe the Aufbau principle | rule that tells how electrons fill orbitals in order of increasing energy; e- occupy lowest energy available orbitals first |
| Describe Hund’s rule | one electron goes into each orbital before any orbital gets two and all single direction electrons have the same spin direction |
| Why is the underlying reason Hund’s rule is followed? | electrons stay unpaired as long as possible to minimize repulsion |
| In general, how do electrons fill atomic orbitals in transition metals? | transition metals fill 4s before 3d, spread e- singly across the five d orbitals (Hund's rule), and lose 4s e- first when forming ions |
| What is the periodic law? | when elements are arranged in order of increasing atomic number, their physical and chemical properties show a repeating pattern |
| Distinguish valence electrons and core electrons | valence- electrons in the outermost shell of an atom core- all other electrons that are not valence electrons |
| Use radial probability distribution functions to show that the core electrons in sodium are on average closer to the nucleus than valence electrons | radial probability distribution shows that sodium's 1s, 2s, and 2p e- have peaks very close to the nucleus while the 3s valence e- peak further outside |
| The number of valence electrons for all the elements in a group is the same (true/false) | true |
| Valence electrons are not involved in the formation of chemical bonds (true/false) | false |
| Valence electrons in main group elements are in the outermost principal level, and therefore are all described by the same principal quantum number (true/false) | true |
| An atom’s valence electrons all have the same angular momentum quantum number (true/false) | false |
| An atom’s valence electrons all belong to the same subshell (true/false) | false |
| An atom’s core electrons are always part of the same principal shell (true/false) | false |
| Arsenic (As) has a total of 18 core electrons (true/false) | false |
| An atom’s d electrons are part of the core, once the outermost d subshell is filled (true/false) | true |
| All halogens have an ns2np5 valence electron configuration (true/false) | true |
| The main-group elements include only those elements found in the s and p x blocks of the periodic table (true/false) | true |
| The row number of an element is equal to the highest principal quantum number of an electron in a ground state atom of that element (true/false) | true |
| Valence electrons are the most tightly held electrons in an atom (true/false) | false |
| What is the trend in effective nuclear charge across a period? | increases across a period from left to right |
| What is the trend in effective nuclear charge down a group? | slightly decreases down a group b/c of increased shielding and distance reduces the nucleus's pull on valance e- |
| Why are differences in effective nuclear charge smaller between consecutive atoms moving down a group compared to moving across the period? | because increases in nuclear charge are nearly canceled by large increases in shielding and distance, whereas across a period shielding stays constant and nuclear charge rises sharply |
| The effective nuclear charge is greater than the nuclear charge (t/f) | false |
| Core electrons shield outer electrons better than outer electrons shield other outer electrons (t/f) | true |
| The effective nuclear charge increases across a period from left to right (t/f) | true |
| The change in effective nuclear charge in a group is typically more pronounced than in a period (t/f) | false |
| How is atomic radius defined? | distance from the nucleus to the outermost e- shell of an atom |
| Metallic radii, nonbonding atomic radii, and covalent radii | metallic- 1/2 the distance b/w the nuclei of two adjacent metal atoms in a metallic lattice; used in metals |
| What is the trend in atomic radius across a group? | Increase b/c as you move across, each element gained an additional e- shell, placing valence e- farther from the nucleus and increasing atomic size. The added inner shells create strong shielding, which reduced effective nuclear charge felt by outer e- |
| What is the trend in atomic radius down a period? | Decreases b/c each element gains a proton while adding e- to the same E level, so shielding doesn't increase. The effective nuclear charge then rises, pulling valence e- closer to the nucleus. This causes the atomic radius to decrease. |
| Covalent radii | covalent- 1/2 the distance b/w the nuclei of two atoms that are covalently bonded |
| Nonbonding radii | nonbonding atomic- 1/2 the distance b/w the nuclei of two atoms that are not bonded but are in close proximity |
| Ionization energy | amount of energy required to remove an electron from an isolated gaseous atom or ion |
| Write the thermochemical equations for the removal of both valence electrons of magnesium. Which process has the highest ionization energy? | First ionization: Mg(g)→Mg+(𝑔)+𝑒 Second ionization Mg+(𝑔)→Mg2+(𝑔)+𝑒− The second ionization has the higher energy b/c the e- is being removed from a positively charged ion where the remaining e- feel a stronger effective nuclear charge |
| What are the general trends for ionization energy in the main group? | increases across a group and decreases down a period |
| Elements in the boron group have ionization energies that are lower than expected. Explain | the first valence e- in Group 13 is in the p orbital while the element to the left (Group 2) has its valence e- in an s orbital; p orbitals are higher in energy and less penetrating than s orbitals |
| Elements in the oxygen group have ionization energies that are lower than expected. Explain | oxygen is in the p subshell, ad one of the p orbitals contains a pair of electrons; paired electrons repel each other, raising their energy and making one easier to remove |
| Which period 3 elements has the largest second ionization energy? | Sodium has the largest second ionization energy. The removal of a second electron from sodium breaks the noble gas core. |
| Electron affinity | energy absorbed or released when an electron is added to a gas phase atom or ion |
| Write the thermochemical equations for the addition of two electrons to an oxygen atom. Which process has the highest (most positive) electron affinity? | O(g) + e– → O–(g) O–(g) + e– → O2–(g) second electron affinities are always positive due to the repulsions and inherent increase in potential energy when an electron is added to an ion that already has a negative charge |
| First electron affinities can be either endothermic (+) or exothermic (–) | first e- affinity: (exothermic) e- is added and configuration becomes more stable, (endothermic) adding e- makes the configuration less stable for elements like noble gases and energy has to be absorbed to overcome repulsive forces |
| Metals form positive ions | metals tend to lose e- to achieve a more stable configuration resulting in cations; metals have low ionization energy, making it easy to give up outer e- |
| Nonmetals form negative ions | nonmetals tend to gain e- to complete their valence shells resulting in anions; nonmetals have large ionization energies so they are likely to accept e- |
| Noble gases do not readily form ions | noble gases have full valence shells making them very stable, and they have little tendency to gain or lose e- |
| What are the characteristics of metals? | good conductors of heat + electricity, solids at room temp, shiny, malleable, ductile, tend to lose e- to form cations |
| What are the characteristics of nonmetals? | poor conductors of heat and electricity, exist in a variety of states and colors, tend to gain e- to form anions |
| What are the trends in metallic character across the periodic table? | metals lose e- and have low ionization energies, ionization increases left to right, elements on the right have high ionization energies, metallic character increases right to left |
| Second electron affinities are always endothermic (+). Explain | adding an e- to a negatively charged ion results in increased repulsion b/w forces, making it harder to absorb a second e-, so energy must be absorbed to overcome this repulsion |
| When atom loses electrons to form an ion, which electrons are lost? | e- from the outermost shell are lost |
| To which orbital are electrons added when a main group atom gains an electron to form an anion? | lowest energy orbital that has a vacancy |
| What do you call particles that have the exact same electron configuration? | isoelectronic |
| Elements tend to lose or gain electrons to have the same number of electrons as: | their nearest noble gas |
| An ion can form by either the addition or removal of protons or by the addition or removal of electrons (T/F) | false |
| When electrons are added to an atom to form anions, the electrons are placed in the largest, highest energy orbitals available (T/F) | false |
| Atoms generally form stable ions when they lose or gain enough electrons so that they have the same number of electrons as their nearest noble gas (T/F) | true |
| Noble gases have especially stable electron configurations in part because of their high ionization energies and endothermic electron affinities (T/F) | true |
| Cations form when the electrons closest to the nucleus are removed from an atom (T/F) | false |
| When orbitals become stabilized, their energy is lowered. When orbitals become destabilized, their energy is raised (T/F) | true |
| What happens to the energies of orbitals as they are filled with electrons as one move across a period? What are the consequences? | energies of orbital decrease as they are filled with e- and orbitals become more stable; atomic size decreases, IE increases, EA becomes more exothermic, chemical reactivity changes |
| In the formation of a transition metal ion, d electrons are removed before s electrons because they are closer to the nucleus (t/f) | false |
| Electrons in the outermost d orbitals lie, on average, closer to the nucleus than electrons in the outermost s orbitals (t/f) | true |
| As transition metal atoms are built up by adding more protons and electrons, the s orbitals stabilize faster (become lower in energy) than the d orbitals (t/f) | false |
| When is an electron configuration said to be paramagnetic? | when it contains unpaired e- in its orbitals |
| When is an electron configuration said to be diamagnetic? | when all e- are paired in its orbitals |
| Define ionic radius | measure of the size of an ion in a crystal lattice, typically defined as the distance from the nucleus of the ion to the outermost electron in its electron cloud; cations have smaller ionic radii than anions |
| Ionic radii of cations are ____________________ than their parent atoms | smaller |
| Ionic radii of anions are ____________________ than their parent atoms | larger |
| Ionic radii get ____________________ as ions become more positively charged. Explain | smaller b/c effective nuclear charge increases, pulling stronger on the remaining e-, and electron-electron repulsions are decreased allowing the nucleus to pull the e- inward |
| Ionic radii get ____________________ as you progress across a group. Explain | larger b/c more electron shells are added which increases the overall size of an ion and distance from the nucleus increases |
| What is an isoelectronic series? | group of ions or atoms that have the same number of electrons but different nuclear charges; more protons = smaller ionic radius, less protons = large ionic radius |
| What is an ionic bond? | type of chemical bond formed through the electrostatic attraction between oppositely charged ions |
| What is an ionic compound? | chemical compound composed of positive ions (cations) and negative ions (anions) held together by ionic bonds |
| What is the energy version of Coulomb’s law? | 𝐸 = 𝑘𝑒 ( 𝑞1𝑞2/𝑟) |
| The potential energy between two oppositely charged particles is positive. (t/f) | false |
| Attractive forces stabilize oppositely charged particles. (t/f) | true |
| As the charges of the cations and anions in a crystal lattice increase, the potential energy becomes less negative. (t/f) | false |
| As the distance between two particles with opposite charges decreases, the potential energy becomes more negative. (t/f) | true |
| Ionic bonds form when oppositely charged ions are attracted to each other by electrostatic forces. (t/f) | true |
| Binary compounds comprised of a metal and nonmetal are usually considered to be ionic. (t/f) | true |
| What is a Lewis-dot symbol? | representation of an atom that shows its valence electrons as dots around the element's chemical symbol |
| What is an octet? | the stable arrangement of eight electrons in the outermost shell (valence shell) of an atom |
| What is the octet rule? | states that atoms tend to gain, lose, or share electrons in chemical bonding to achieve a stable configuration with eight valence electrons in their outermost shell |
| What are the general properties of ionic compounds? | Have high melting points, high boiling points, tend to be hard and brittle, tend to be good insulators in the solid state, conduct electricity when molten, conduct electricity when dissolved in water |
| Why don’t ionic compounds conduct electricity in the solid state? Can ionic compounds conduct electricity? If so, how? | when solid, ions are fixed, preventing electrical conductivity; when ions are dissolved, they can conduct electricity b/c their ions are mobile |
| Ionic substances have low enthalpies of fusion and therefore melt relatively easily. (t/f) | false |
| Compounds that contain ionic bonds tend to fracture if enough force is applied to the compound. (t/f) | true |
| An ionic compound always conducts electricity because it contains ions. (t/f) | false |
| The electrostatic forces between the ions in an ionic solid are stronger than the electrostatic forces between the molecules in a molecular solid. (t/f) | true |
| What is a crystal lattice? | 3-D ordered arrangement of ions in an ionic solid where each cation is surrounded by a anions and vice verse, held together by electrostatic attractions |
| What is lattice energy? | energy released when gaseous ions come together to form one mole of an ionic solid |
| Lattice energy is negative. When the lattice energy of two substances is compared, which substance has the stronger ionic bond (the one with the most negative or the least negative lattice energy)? | the most negative lattice energy corresponds to the stronger ionic bond |
| What is the relationship between the separation of ions and the lattice energy? | as the separation of the ions increases (r), the lattice energy decreases (becomes less negative) |
| What is the relationship between the product of ionic charges and the lattice energy? | as the product of the charges increases, the lattice energy increases (becomes more negative) |
| The more negative the lattice energy, the weaker the ionic bond. (t/f) | false |
| Lattice energy is the amount of energy needed to precipitate a salt from an aqueous solution of its ions. (t/f) | false |
| The potential energy between oppositely charged ions becomes more negative as their internuclear separation decreases. (t/f) | true |
| As the charge on either the cation or anion increases, the lattice energy becomes less negative. (t/f) | false |
| Charge has a bigger impact on lattice energy than the internuclear distance between ions. (t/f) | true |
| What is a Born-Haber cycle? | thermochemical process used to calculate the lattice energy of an ionic compound |
| A Born-Haber cycle equates two pathways, path 1 and path 2. What is true about the total energies of path 1 and path 2? | path 1- direct formation of the ionic compound from its elements in their standard states (enthalpy of formation) path 2- sum of all intermediate steps, including sublimation, ionization, dissociation, electron affinity, and lattice energy |
| One of the steps in the pathways of the Born-Haber cycle involves breaking covalent bonds to form atoms from a gas phase molecule. What is the energy associated with this process called? | bond dissociation energy (or bond enthalpy) |
| What is a covalent bond? | occurs when 2 atoms share 2 or more pairs of e- |
| What is a covalent compound? | one that contains only covalent bonds |
| What is a molecular compound? Are covalent and molecular compounds necessarily the same thing? | molecular compound is a chemical composed of molecules that are bonded together by covalent bonds; covalent compounds are held together by covalent bonds whole molecular compound are made of discrete, individual molecules |
| What is bond length? | intranuclear distance between two bonded atoms |
| Atoms in covalent bonds share electrons (t/f) | true |
| Molecular compounds tend to have high melting points (t/f) | false |
| Molecular compounds tend to be more flammable than ionic compounds (t/f) | true |
| At room temperature, it is common to find examples of covalent compounds in the solid, liquid and gas phase (t/f) | true |
| Covalent bond formation is a destabilizing force and causes potential energy to increase (t/f) | false |
| The bond length of a covalent bond is defined as the internuclear distance between the nuclei of the atoms involved in the bond (t/f) | true |
| Covalent bond formation is an exothermic process (t/f) | true |
| In covalent bonding, the repulsive forces between the nuclei are overcome by the attractive forces between the electrons of one atom and the nuclei of the other (t/f) | true |
| Distinguish single, double, and triple bonds | single: 1 bond pair, weak, long double: 2 bond pairs, intermediate strength, intermediate length triple: 3 bond pairs, strongest, shortest |
| Nonpolar covalent bonds | nonpolar covalent: e- are shared equally b/w atoms, small electronegativity difference, bonds are in diatomic molecules |
| What is a dipole? | a situation where there is a separation in electrical charges within a molecule, leading to one end having a partial + and the other having a partial - |
| What is electronegativity? What is the general trend in electronegativity across the periodic table? | electronegativity is the measure of an atoms ability to attract and hold shared electron in a chemical bond: increase across a group and decreases down a period |
| How does one determine whether a bond is nonpolar covalent, polar covalent, or ionic? | nonpolar = <0.5 polar covalent = 0.5-2.0 ionic = >2.0 |
| What is a dipole moment? | measure of the separation of + and - charges in a molecule |
| Write the formula for calculating dipole moment and define the terms | μ = q × r (μ is dipole movement, q is charge, r is separation of charges) |
| What is the relationship between dipole moment and polarity? | polarity is the uneven distribution of electrical charge in a molecule and dipole movement quantifies this polarity |
| What is meant by percent ionic character? | measure of how much a chemical bond resembles an ionic bond by comparing the observed dipole movement to the dipole movement it would have if the bond were completely ionic |
| Write the formula for percent ionic character and define the terms | (μ/er) × 100 (μ is the dipole movement, e is the elementary charge, r is the separation between the charges) |
| What is significant about having a bond with a percent ionic character greater than or equal to 50%? | suggests that the bond has substantial ionic character and behaves more like an ionic bond than a covalent bond |
| Are there any bonds with 100% ionic character? | no because quantum numbers overlap, electron density is shared, and there is an electronegativity limit |
| Summarize the process for drawing Lewis-structures for molecular compounds and polyatomic ions | count valence e-, determine central atom, connect outer atoms, place leftover e- around outer atoms in pairs until each non-hydrogen atom has an octet, check for octets, try multiple bonds if needed |
| How does an ionic charge affect the number of valence electrons in a Lewis structure? Comment on the effects of both cations and anions | each negative charge adds one e- to the valence e- count with each positive charges reduces the e- count by 1 |
| Only the valence electrons of a substance are used in the Lewis structure | true |
| When counting valence electrons in polyatomic ions, always add one electron for each positive charge and subtract one electron for each negative charge | false |
| When determining the central atom, the least electronegative atom is usually the central atom, unless the least electronegative atom is hydrogen | true |
| When constructing a Lewis structure, the surrounding atoms are initially connected to the central atom using single covalent bonds. Each bond uses two valence electrons from the total available | true |
| After connecting the surrounding atoms to the central atom, if any additional valence electrons remain, these electrons are placed around the surrounding atoms in pairs, starting with the least electronegative first | false |
| Surrounding atoms are given as many electrons as possible | false |
| After filling the octets of terminal atoms, additional electrons are placed around the central atom in pairs | true |
| If an atom (other than H) does not have an octet, you should try to form multiple bonds | true |
| What are resonance structures? | different ways of drawing the same molecule or ion to represent delocalization of electrons within it |
| What does delocalization of electrons do for a molecule or ion? | increases the stability of a molecule or ion by spreading e- density over multiple atoms or bonds, therefore reducing e- repulsion, lowing overall energy of the system, enhancing stabilty |
| What two concepts should be kept in mind when considering resonance structures? | validity, which states that resonance structures only differ in the position of the e-, and resonance hybrid, which states the actual molecule is not one structure but a hybrid of valid resonance structures |
| What is formal charge? | theoretical charge assigned to an atom in a molecule or ion to help determine the most stable resonance structure |
| How is formal charge calculated? | valence electrons - electrons assigned in Lewis structure (e- assigned in Lewis structure = 1/2 bond pair e- + lone pair e-) |
| How is the best resonance structure—or the largest contributor to the resonance hybrid—identified? | sum of formal charges must = overall charge on molecule, smaller formal charges are preferred over larger formal charges, negative formal charge should reside on the more electronegative atoms |
| Formal charges represent actual charges on atoms | false |
| Formal charges are calculated as the number of valence electrons minus the number of lone pair electrons assigned to an atom in its Lewis structure. | false |
| When determining the best resonance structure, small formal charges are favored over large formal charges. | true |
| The best resonance structure of an anion has a negative formal charge on the most electronegative atom. | true |
| When a molecule can be represented by two resonance structures, the molecule exists as one structure half of the time and the other structure the rest of the time | false |
| All resonance structures must have the same connectivity of atoms; the only difference is the distribution of electrons | true |
| Which elements can expand their octets? | elements in the third period and beyond |
| In situations where a central atom has an expanded octet, what is typically true of the outer atoms. Why is this important? | outer atoms are typically highly electronegative, which likely adds stability by withdrawing a significant amount of e- density away from the central atom |
| When is it common to find molecules with central atoms with less than an octet? | central atoms from groups 2 and 13 because they lack valence e- and such molecules are often reactive and tend to accept e- to achieve a complete octet |
| What is a free radical? | molecule or ion with an unpaired e- in its outer shell making it highly reactive as it seeks to pair the unpaired e- by reacting with other molecules |
| When do molecules or polyatomic ions form free radicals? | when they undergo processes that result in the loss of e-, leading to an unpaired e- in their outermost shell; some process includes redox reactions, heat, electrical discharges) |
| What is bond energy? | energy needed to separate two covalently bonded atoms |
| When a bond is formed is energy absorbed or released? What does this mean for the sign of the energy term? | energy is released, bond energy for bond formation is negative, and bond energy for breaking a bond is endothermin |
| What is bond order? | # of bonds between two atoms |
| Polar covalent bonds | e- are shared unequally, causing partial + and partial -, moderate electronegativity difference |
| Ionic bonds | no sharing of e- but e- are transferred from one atom to another forming ions, large electronegativity difference between |
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