Chemistry Test #2 Chapter 7 and 8
chapter 8: what do do when an ion has a positive or a negative charge? is the process of forming a cation endothermic or exothermic? example: Mg2+ cation what about an anion? is the process of forming an anion endothermic or exothermic? example: S2- anion
cations formed when metal atom loses its valence e-s, resulting in a new lower energy level valence shell. Process is always endothermic. -magnesium atom has 2 valence e-s (Mg atom 1s22s22p63s2) When magnesium forms a cation, loses its valence e-s and assumes noble gas configuration (in this case Ne). Mg2+ cation = 1s22s22p6 -Al atom=1s22s2sp63s63p -Al3+ ion=1s22s22p6 (configuration of neon) -anions formed when nonmetal atoms gain enough electrons to have 8 valence electrons -filling the s and p sublevels of the valence shell -the sulfur atom has 6 valence electrons: S atom=1s22s22p63s23p4 -to have 8 valence electrons, sulfur must gain two more and have electron configuration of Ar: S2- anion=1s22s22p63s23p6
chapter 8: which orbitals are the most stable? how to half filled orbitals compare with other orbitals in terms of stability? exceptions to first ionization energy trends?
full orbitals most stable (noble gases) half filled orbitals more stable than partially filled orbitals Beryllium and boron (boron has smaller IE), aluminum and gallium, nitrogen and oxygen (O has lower first ionization energy)
chapter 7: list all the orbitals in each principal level. specify the three quantum numbers for each orbital a) n=1 b) n=2 c) n=3 d) n=4
n=1 (l values, ml values) (0,0) n=2 (0,0) (1,-1) (1,0)(1,1) n=3 (0,0)(1,-1)(1,0)(1,1)(2,-2)(2,-1)(2,0)(2,1)(2,2) n=4 (0,0)(1,-1)(1,0)(1,1)(2,-2)(2,-1)(2,0)(2,1)(2,2)(3,-3)(3,-2)(3,-1)(3,0)(3,1)(3,2)(3,3)
chapter 7: relationships
n=1+ l=0 --> (n-1) ml=-l to l
chapter 7: Which of the following transitions for an electron in a hydrogen atom would release the largest quantum of energy? n=3 → n=1 n=4 → n=3 n=1 → n=4 n=2 → n=1
n=3 --> n=1
Chapter 8: what does penetration cause with the energies of sublevels in the same principal energy level?
penetration causes energies of sublevels in same principal level to not be degenerate (not have same energy). In 4th and 5th principal levels, effects of penetration become so important that the s orbital lies lower in energy than the d orbitals of the previous principal level
chapter 7: what suggested wave nature for particles? what is spectra? How do you separate light emitted by a single element in a glass tube into its constituent wavelengths? what does this create?
atomic spectroscopy, which is the study of the electromagnetic radiation absorbed and emitted by atoms --> suggested wave nature for particles -when atom absorbs energy, often re emits energy as light. Atoms of each element emit light of a characteristic color. Spectra: when atoms or molecules absorb energy, that energy is often released as light energy (fireworks, neon lights, etc) -Can separate the light emitted by a single element in a glass tube into its constituent wavelengths by passing it through a prism (when that emitted light is passed through a prism, a pattern of particular wavelengths of light is seen that is unique to that type of atom or molecule → pattern=emission spectrum.)
light from 3 different lasers (A, B, and B) each with a different wavelength, was shined onto the same metal surface. laser A produced no photoelectrons. Lasers B and C both produced photoelectrons, but the photoelectrons produced by laser B had a greater velocity than those produced by laser C. arrange the lasers in order of increasing wavelength
B<C<A
for which element is the gaining of an electron most exothermic? a) Li b) N c) F d) B
C
the ionization energies of an unknown third period element are listed here. identify the element: IE1=786 kJ/mol, IE2=1580 kJ/mol, IE3=3230 kJ/mol, IE4=4360 kJ/mol, IE5=16,100 kJ/mol a) Mg b) Al c) Si D) P
C
which electron in sulfur is most shielded from nuclear charge? A) electron in 1s orbital B) electron in 2p orbital C) electron in 3p orbital D) none of the above (all electrons equally shielded from nuclear charge)
C
which wavelength of light has the highest frequency? a) 10 nm b) 10 mm c) 1 nm d) 1 mm
C
arrange these atoms and ions in order of increasing radius: Cs+, Ba2+, I-
Cs+<Ba2+<I-
calculate the wavelength of an electron traveling at 1.85 x 10^7 m/s a) 2.54 x 10^13 m b) 3.93 x 10^-14 m c) 2.54 x 10^10 m d) 3.93 x 10^-11 m
D
which element has the smallest atomic radius? a) C b) Si c) Be d) F
D
which statement is true about effective nuclear charge? a) decreases as move to right across row b) increases as move right across row c) remains constant as move right across row d) increases and decreases at regular intervals as move across row.
D
which statement is true about electron shielding of nuclear charge? a) outermost electrons efficiently shield one another from nuclear charge b) core electrons efficiently shield one another from nuclear charge c) outermost electrons efficiently shield core electrons from nuclear charge d) core electrons efficiently shield outermost electrons from nuclear charge
D
chapter 7: why doesn't low frequency light doesn't eject e-s? what does increasing the frequency of light do to the energy of each photon?
Ejected electrons: -low-frequency light doesn't eject e-s because no single photon has the minimum energy necessary to dislodge the e- -increasing the frequency of light (even at low intensity) increases the energy of each photon -one photon at the threshold frequency gives the electron just enough energy for it to escape the atom (binding energy) -when irradiated with a shorter wavelength (higher frequency) photon, the electron absorbs more energy that is necessary to escape. This excess energy becomes KE of the ejected electron. KE=E(photon)-E(binding)
Chapter 8: what is coloumb's law? for like charges, what is the potential Energy? what about for opposite charges? how does the strength of the interaction change as the size of the charges increases?
The only forces acting in an atom are coulomb forces. -coulomb's law describes the attractions and repulsions between charged particles -for like charges, the potential energy E is positive and decreases as the particles get farther apart (as r increases.) → like charges repel! -for opposite charges, the potential energy is negative and becomes more negative as the particles get closer together (r decreases) → opposite charges attract! -the strength of the interaction increases as the size of the charges increases (e-s more strongly attracted to a nucleus with a 2+ charge than a nucleus with a 1+ charge)
chapter 7: p orbitals
I=1, p orbitals: px, py, and pz orbitals. Each principal level with n=2 or greater contains 3 p orbitals -each principal energy state above n=1 has 2 p orbitals (mi=-1,0,+1) -each of 3 orbitals points along different axis (px, py, pz) -second lowest energy orbitals in principal energy state, two-lobed
chapter 7: d orbitals?
I=2, d orbitals. Each principal level with n=3 or greater contains 5 d orbitals (ml=-2 → 2) -each principal energy state above n=2 has 5 d orbitals (ml=-2,-1,0,+1,+2) -⅘ orbitals aligned in a different plane. 5th aligned w z axis, dz squared -dxy, dyz, dxz, dxsquared - y squared -third lowest energy orbitals in principal energy level, mainly 4-lobed (but 1 two-lobed w toroid) -planar nodes (higher principle levels also have spherical nodes)
Chapter 8: what is penetration? as an outer electron undergoes penetration into the region occupied by inner electrons, what does it experience? what is the degree of penetration related to?
Penetration: the closer an electron is to the nucleus, the more attraction is experiences. The better an outer electron is at penetrating through the electron cloud of inner electrons, the more attraction it will have for the nucleus (the lower its energy will be) -as an outer electron undergoes penetration into the region occupied by inner electrons, experiences a greater nuclear charge and therefore a lower energy! -the degree of penetration is related to the orbital's radial distribution function. In particular, the distance the maxima of the function are from the nucleus
chapter 7: schrodinger's equation what does the plot of the wave function squared represent what are the 3 integer terms in the wave function that are its solutions?
Schrodinger's equation Allows us to calculate the probability of finding an electron with a particular amount of energy at a particular location in the atom. Solutions to the equation produce many wave functions (w) -plot of wave function squared represents an orbital, a position probability distribution map of the e- Solutions to the wave function: -size, shape, and orientation in space of an orbital are determined to be 3 integer terms in the wave function (added to quantize the energy of the electron) -these integers=quantum numbers (principal quantum number=n, angular momentum quantum number=I, magnetic quantum number=mI)
chapter 7: why are atoms drawn as spheres?
Shape of atoms: atoms drawn as spheres because most atoms contain many e-s occupying a number of different orbitals. Therefore, the shape of an atom is obtained by superimposing all of its orbitals (if we superimpose the s, p, and d orbitals, get a roughly spherical shape)
arrange these elements in order of increasing first ionization energy: Cl, Sn, Si
Si<Cl<Sn
chapter 7: what is the problem with the photoelectric effect? what is a threshold frequency? does low-frequency light eject electrons from a metal? does high frequency light eject electrons from a metal?
-a min frequency was needed before electrons would be emitted regardless of the intensity=threshold frequency -high-intensity light from a dim source caused electron emission without any lag time -light used to dislodge e-s in photoelectric effect exhibits a threshold frequency, where no e-s are emitted from the metal no matter how long the light shines on it. -low-frequency light does no eject electrons from a metal regardless of its intensity or its duration -high-frequency light does eject electrons.
chapter 8: what is paramagnetism? when does this occur? what is diamagnetism? when does this occur?
-e- configurations that result in unpaired e-s mean that the atom or ion will have a net magnetic field → paramagnetism (will be attracted to a magnetic field) (if one sublevel only slightly filled, paramagnetic) -e- configurations that result in all paired e-s mean that the atom or ion will have no magnetic field → diamagnetism (slightly repelled by magnetic field) (if all sublevels filled, diamagnetic)
Chapter 8: shielding and effective nuclear charge in multi-electron atoms, how does shielding relate to an electron's attraction to the nucleus and repulsion by other electrons in the atom? what is the effective nuclear charge experienced by an electron?
-each electron in a multi-electron atom experiences both attraction to nucleus and repulsion by other electrons in the atom. These repulsions cause e- to have a net reduced attraction to the nucleus; e- shielded from nucleus/full effects of nuclear charge -Repulsion of 1 electron by other electrons=shielding (shielding e- from full effects of nuclear charge) -total amount of attraction that e- feels for nucleus=effective nuclear charge experienced by e-
chapter 8: what are the two factors that the strength of attraction is related to? if an electron is in a larger orbital in comparison to another, will it have more attraction or less attraction to the nucleus?
-strength of attraction related to most probable distance the valence e-s are from nucleus and the effective nuclear charge the valence electrons experience -larger the orbital an electron is in, farther its most probable distance will be from the nucleus and the less attraction it will have for the nucleus. -traversing across a period increases the effective nuclear charge on the valence electrons (penetration)
chapter 8: what are the two ways to define the size of an atom? how does this differ for nonmetals and metals in a covalent radius? what is the definition of the atomic radius?
-two ways to define the size of an atom -van der waals radius=nonbonding -covalent radius=bonding atomic radius -nonmetals: ½ the distance between 2 of atoms bonded together -metals: ½ of distance between 2 atoms next to each other in a crystal of the metal. -atomic radius is an average radius of an atom based on measuring large numbers of elements and compounds. Refers to set of average bonding radii determined from measurements on a larger number of elements and compounds. Represents radius of an atom when bonded to another atom. Always smaller than van der waals radius.
Chapter 8: how many valence electrons and core electrons are in Kr? how does the number of valence electrons relate to the periodic table? how does the principal energy level of the valence electrons relate to the periodic table?
28 core, 8 valence. to find number of valence electrons, look at group number! period corresponds to principal energy level of valence electrons
how much time (in seconds) does it take light in a vacuum to travel 1.00 billion km?
3.33 x 10^3 s
specific heat capacity of water
4.186
according to colomb's law, if the separation between 2 particles of the same charge is doubled, what happens to the potential energy of the two particles? A) it is twice as high as it was before the distance separation B) it is 1/2 as high as it was before the separation C) it does not change D) it is 1/4 as high as it was before the separation
B
chapter 7: Heisenberg uncertainty principle why can an electron be both a particle and a wave?
Uncertainty principle: Heisenberg stated that the product of the uncertainties in both the position and speed of a particle was inversely proportional to its mass: know equation! -more accurately you know the position of a small particle, such as an electron, the less you know about its speed, and vice versa. -electron can be both a particle and a wave because of complementarity
electron configuration for Fe2+?
[Ar]4s^03d^6
chapter 7: determinacy vs indeterminacy Are newton's laws deterministic?
-according to classical physics, particles move in a trajectory/path determined by the particle's velocity, positions, and the forces acting on it. -determinacy=definite, predictable future. -because we cannot know both the position and velocity of an electron, we cannot predict the path it will follow -Newton's laws deterministic: past determines future. -indeterminacy=indefinite future, can only predict probability. Future path of an e- =indeterminate, can only be described statistically. The best we can do is to describe the probability an electron will be found in a particular region using statistical functions and prob distribution maps (statistical map that shows were e- likely to be found under given set of conditions)
Chapter 8: how do energy levels and sublevels fill? what is the aufbau principle? how is the energy related for orbitals in the same sublevel? hund's rule?
Filling the orbitals with electrons: energy levels and sublevels fill from lowest energy to high: aufbau (building up) principle → building up electron configurations for elements. -orbitals that are in the same sublevel have the same energy (degenerate) -no more than 2 electrons per orbital: pauli exclusion principle -when filling orbitals that have the same energy, place 1 e- in each before completing pairs: hund's rule → when filling degenerate orbitals, electrons fill them singly first, with parallel spins. E-s in all sublevels with the highest principal energy shell=valence electrons. Electrons in lower energy shells=core electrons. Summary: -e-s occupy orbitals to minimize the energy of the atom; lower energy levels fill first. -orbitals can hold no more than 2 e-s each. When two e-s occupy orbital, spins opposite. -when orbitals of identical energy available, e-s occupy orbitals singly with parallel spins rather than in pairs.
chapter 8: halogens In their reactions with metals, do they tend to gain or lose electrons? what about in their reactions with nonmetals?
Halogens: most reactive nonmetals. have 1 fewer e- than next noble gas. In their reactions with metals, halogens tend to gain e- and attain e- configuration of the next noble gas, forming anion with charge 1-. -in their reactions with nonmetals, tend to share e-s with other nonmetal so that each attains the e- configuration of a noble gas
chapter 8: trends in affective nuclear charge? what is the trend in atomic radius as you move across a row determined by?
-as go down group, nuclear charge decreases -as go across period, nuclear charge increases -trend in atomic radius as move to right across a row in periodic table determined by inward pull of nucleus on electrons in outermost principal energy level. -any one e- in a multielectron atom experiences both the + charge of the nucleus and the - charges of other electrons -core electrons efficiently shield electrons in the outermost principal energy level from nuclear charge, but outermost electrons do not efficiently shield one another from nuclear charge -as move down column, n of electrons in outermost principal energy level increases, resulting in larger orbitals and larger atomic radii -as move right across row, effective nuclear charge experienced by e-s in outermost principal energy level increases, resulting in stronger attraction between outermost e-s and nucleus, and smaller atomic radii
chapter 8: with main group elements, how does atomic radius change as you go across a period? what about as you go down a group/column? Arrange in order of decreasing radius: S, Ca, F, Rb, Si
-atomic radius decreases across period (left to right). Adding electrons to same valence shell, effective nuclear charge increases. Valence shell held closer -atomic radius increases down the group/column: valence shell farther from nucleus. Effective nuclear charge fairly close because of effective shielding. Rb, Ca, Si, S, F
chapter 8: alkali metals what are the periodic trends for ionization energy and atomic radius of alkali metals? what about for e- affinity and melting point? density? are there ionization energies high or low? does the increase in mass outplace the increase in volume, or does the increase in volume outpace the increase in mass? what determines if an alkali metal is more reactive?
-atomic radius increases down column, ionization energy decreases down column (and increases across period) -very low ionization energies: easier to oxidize so good reducing agents, very reactive and not found uncombined in nature, react with nonmetals to form salts, compounds generally soluble in water (found in seawater) -e- affinity decreases down column, melting point decreases down column (except K) -in general, increase in mass > increase in volume (outpaces increase in volume) caused by greater atomic radii as move down column -density increases as move down column w exception of K -lower the first ionization energy of alkali metal, greater tendency it will have to lose its e- and the more reactive it is.
chapter 7: Bohr model of the atom? why was the atom model developed? major idea of his model?
-developed to explain how the structure of the atom changes when it undergoes energy transitions. Major idea=energy of the atom was quantized (the amount of energy in the atom was related to the electron's position of the atom.) -quantized=the atom could only have very specific amounts of energy The energy of the electron is proportional to the distance the orbit is from the nucleus Bohr: attempted to find a model for the atom that explained atomic spectra. E-s travel around nucleus in circular orbits that existed at specific, fixed distances from nucleus. Energy each orbit=quantized. Orbits called "Stationary states" and possessed a stability. Also thought that radiation is only emitted/absorbed when electron jumps/transitions from one stationary state to another -emitted radiation=photon of light distance between orbits determined energy of photon of light produced.
chapter 7: what did Einstein propose about light?
-einstein proposed that the light energy was delivered to the atoms in packets called quanta or photons (packet of light=photon or quantum) -the energy of a photon of light is directly proportional to its frequency -inversely proportional to its wavelength -proportionality constant=planck's constant (h), value 6.626x10-34 J x s -emission of electrons from the metal depends on whether or not a single photon has sufficient energy (as given by hv) to dislodge a single electron
chapter 7: is the emission spectrum of a particular element always the same or different? is white light spectrum continuous or discontinuous? what can spectra be used for?
-emission spectrum of a particular element is always the same, so can use emission spectrum to identify the element. -white light spectrum continuous, while some are discontinuous (bright lines at specific wavelengths with complete darkness in between) -if noncontinuous, can identify the material through flame tests (Na=orange flame, K=purple flame, Li=red flame, Ba=white flame) Spectra: can be used to identify atoms and molecules. Like unique fingerprints. Instrument used to measure spectra=spectrometer.
chapter 7: what does an emission spectrum for an atom look like? What is the wave behavior of electrons, according to de Broglie?
-emission spectrum of atom that contains discrete lines: stationary states only exist at specific, fixed energies. Wave behavior of electrons (Broglie): de Broglie proposed that particles could have wavelike character. Predicted that the wavelength of a particle was inversely proportional to its momentum. Because the electron is so small, the wave character of electrons is significant.
chapter 7: interference and diffraction constructive interference vs destructive?
-interaction between waves=interference. When waves cancel each other out or build each other up depending on their alignment upon interaction. -constructive interference: if two waves of = amplitude in phase when they interact (align with overlapping crests), a wave with twice the amplitude results. -destructive interference: if 2 waves out of phase (align so crest of one overlaps with trough of other) when interact, waves cancel each other by destructive interference. Diffraction: -when traveling waves encounter an obstacle or slit in a barrier that is about the same size as the wavelength, they bend/diffract around it; called "diffraction". Traveling particles do not diffract
chapter 8: what is ionization energy? what is the first ionization energy? second ionization energy? the larger effective nuclear charge on an electron, the.... the farther the most probable distance an e- is from the nucleus, the.... Ionization trends on periodic table?
-ionization energy of atom/ion =the min energy required to remove an electron from the atom or ion in the gaseous state. Always + -first ionization energy=energy to remove first e- from neutral atom, second IE=energy to remove 2nd electron from 1+ ion, etc. -larger effective nuclear charge on the e-, more energy takes to remove it. -farther the most probable distance e- is from nucleus, less energy it takes to remove it. First IE generally decreases down column/family because e-s in outermost principal level are increasingly farther away from + charged nuclear and are therefore held less tightly. (valence e- farther from nucleus) -first IE generally increases across period bc e-s in outermost principal energy level generally experience a greater effective nuclear charge
Chapter 8: what does mendeleev's periodic law allow us to predict? how is the periodic table organized? what does quantum mechanics explain? what is a periodic property?
-mendeleev's periodic law allows us to predict what the properties of an element will be based on its position on the table (doesn't explain why pattern exists) -table organized according to periodic law: when elements arranged in order of increasing mass, certain properties recur periodically. Mass increases from left to right, elements w similar properties fall in same column -quantum mechanics= Explains the underlying reasons for the periodic table, explains how e-s arranged in atoms, which in turn determines the element's properties. -periodic property: one that is predictable based on an element's position within the periodic table. -laws summarize behavior, theories explain behavior.
Chapter 7: what do newton's laws predict? does the particle nature or wave nature of electrons dominate? when confined to a very small volume in atoms, what properties to electrons exhibit?
-newton's laws (F=ma) predict the exact paths of electrons. Electrons follow predictable paths when they are not confined to small volumes. Their particle nature dominates. -electric field exists between the plates exerting a force on the negatively charged electron The strange behavior of particles with tiny masses confined to very small volumes: -electrons are incredibly small in mass, confined to a very small volume in atoms, when residing within atoms, e-s exhibit both a particle nature(mass) and wave nature(wavelength) -when unobserved, absolutely small particles like electrons can be in 2 different states at the same time.
chapter 7: what do nodes represent in regards to probability?
-nodes in functions are where prob drops to 0. A point where the wave function, and therefore the probability density and radial distribution function, all pass through 0. -net result=plot that indicates most probable distance of e- in a 1 s orbital of H is 52.9 pm=52.9x10-12 m=52.9x10-10 cm=0.529x10-8cm=0.529 Angstrom
chapter 8: how do the radii of transition elements differ across each row?
-radii of transition elements stay roughly constant across each row. -atoms in the same group increase the size down the column. -Atomic radii of transition metals are roughly the same size across the d block
Chapter 8: do s orbitals penetrate more fully or less than p orbitals do? do p orbitals penetrate more fully than d orbitals do? etc. how are the affects of penetration seen as important within the 4th and 5th principal levels?
-s orbitals penetrate more fully than p orbitals, which penetrate more fully than d orbitals, etc. -because of penetration, sublevels of each principal level not degenerate for multielectron atoms -4th and 5th principal levels, effects of penetration become so imp that 4s orbital lies lower in energy than 3d orbitals and 5s orbital lies lower in energy than 4d orbitals -energy separations between 1 set of orbitals and the next become smaller for 4s orbitals and beyond.
Chapter 8: review this!
-s sublevel has 1 orbital: can hold 2 e-s -p sublevel has 3 orbitals: can hold 6 e-s -d sublevel has 5 orbitals; can hold 10 e-s -f sublevel has 7 orbitals; can hold 14 e-s
Chapter 8: what does the splitting of sublevels within a principal level lead to? electrons in a 2p orbital have a greater/lesser probability of being found closer to the nucleus than an electron in a 2s orbital
-splitting of sublevels within principal level result of spatial distributions of electrons within a sublevel. -E- in a 2p orbital has greater probability of being found closer to the nucleus than an electron in a 2s orbital (2s orbital is lower in energy only when the 1s orbital is occupied.) -most of the probability in a radial distribution function of the 2p orbital lies outside the radial distribution function of the 1s orbital, so almost all of the 2p orbital is shielded from nuclear charge by the 1s orbital → result=2s orbital, since experiences more of nuclear charge due to its greater penetration, is lower in energy than 2p orbital.
chapter 8: FIGURE THIS OUT
-the larger the effective nuclear charge an e- experiences, the stronger the attraction it will have for the nucleus -the stronger the attraction the valence e-s have for the nucleus, the closer their average distance will be to the nucleus -traversing across a period increases the effective nuclear charge on the valence electrons. down group, nuclear charge decreases. across period, nuclear charge increases.
chapter 7: what is a probability density function ? where does it have a maximum? does it increase or decrease with increasing r? at the nucleus, what is the probability density? what about the radial distribution function?
-the probability density function, which is the probability per unit volume, has a maximum at the nucleus and decreases with increasing r -volume of the thin shell, which is 0 at the nucleus and increases with increasing r -at nucleus, probability density is at a maximum, but volume of shell is 0, so radial distr function=0. As r increases, volume of thin spherical shell increases
Chapter 8: does the 2s orbital or 2p orbital penetrate more deeply into the 1s orbital? what does a weaker penetration entail for the electrons in a sublevel? what does a deeper penetration entail for the electrons in a sublevel? the lower the value of the l quantum number, the less....
-the radial distribution function shows that the 2s orbital penetrates more deeply into the 1s orbital that does the 2p. -weaker penetration of 2p sublevel → electrons in 2p sublevel experience more repulsive force; more shielded from attractive force of nucleus -deeper penetration of 2s electrons → e-s in 2s sublevel experience greater attractive force to nucleus, not shielded as effectively -the lower the value of the l quantum number, the less energy the sublevel has: s (l=0)<p(l=1)<d(l=2)<f(l=3)
chapter 8: the larger the orbital an electron is in, the farther.... the larger the principle energy level an orbital is in, the larger...
-the size of an atom is related to the distance the valence electrons are from the nucleus. The larger the orbital an electron is in, the farther its most probable distance will be from nucleus, and the less attraction it will have for the nucleus -traversing down a group adds a principal energy level -the larger the principal energy level an orbital is in, the larger its volume
chapter 8: how to find the charge of an ion that metals and nonmetals form based on looking at the periodic table. what is the electron configuration like for the atoms that form ions?
-we have seen that many metals and nonmetals form 1 ion. In some, the charge on that ion is predictable based on its position on the periodic table (Group 1A=1+, group 2A=2+, group 7A=1-, group 6A=2- etc.) These atoms form ions that will result in an e- configuration that is the same as the nearest noble gas.
Chapter 8: as you move down a column, how does the number of electrons in the outermost principal energy level change? what are valence electrons? for main-group elements, where are they located? what are core electrons?
-with each subsequent row, highest principal quantum number increases by 1 -as move down a column, the number of electrons in the outermost principal energy level (highest n value) remains the same -electrons in all sublevels with highest principal energy shell=valence electrons -for main-group elements, valence e-s are those in outermost principal energy level. Electrons in lower energy shells=core electrons. One of most imp factors in the way an atom behaves, both chemically and physically, is the number of valence electrons.
calculate the wavelength of light emitted when an electron in the hydrogen atom makes a transition from an orbital with n=5 to an orbital with n=3
A
which electron transition produces light of the highest frequency in the hydrogen atom? a) 5p--> 1s b) 4p --> 1s c) 3p --> 1s d) 2p --> 1s
A
which set of 4 quantum numbers corresponds to an electron in a 4p orbital? a) n=4, l=1, ml=0, ms=1/2 b) n=4, l=3, ml=3, ms=-1/2 c) n=4, l=2, ml=0, ms=1/2 d) n=4, l=4, ml=3, ms=-1/2
A
chapter 8: alkali metals in their reactions, do they tend to gain or lose electrons?
Alkali metals: most reactive metals. have 1 more e- than previous noble gas. In their reactions, alkali metals tend to lose 1 e-, resulting in the same electron configuration as a noble gas (forming cation with 1+ charge) -elements w e- configurations close to those of noble gases are most reactive because can attain noble gas e- configurations by losing or gaining a small number of electrons.
chapter 7: amplitude of a wave vs wavelength? how do different amplitudes and wavelengths affect light? what is the frequency of a wave? what is it proportional to? inversely proportional to? what is the total energy proportional to?
Amplitude: height of the wave/vertical height of crest. Distance from node to crest or node to trough. amplitude=measure of light intensity (larger amplitude, brighter light) Wavelength: distance between adjacent crests, measure of the distance covered by the wave. Distance from 1 crest to next or distance between alternative nodes. Different wavelengths produce different colors of lights, different amplitudes produce different brightnesses of light. frequency(v) is the number of waves/cycles that pass through a stationary point in a given period of time. Frequency proportional to speed at which wave is traveling. inversely proportional to wavelength. Total energy is proportional to the amplitude of the waves. The larger the amplitude, the more force it has.
how much energy (in J) is contained in 1 mol of 552 nm-photons? a) 3.6 x 10^-19 J b) 2.17 x 10^5 J c) 3.60 x 10^-28 J d) 5.98 x 10^-43 J
B
which kind of electromagnetic radiation contains the greatest energy per photon?
B
which set of 3 quantum numbers does not specify an orbital in the hydrogen atom? a) n=2, l=1, ml=-1 b) n=3, l=3, ml=-2 c) n=2, l=0, ml=0 d) n=3, l=2, ml=2
B
which species is diamagnetic? a) Cr2+ b) Zn c) Mn d) C
B
which statement is true about trends in metallic character? a) metallix character increases as move to right across row and increases as move down column b) decreases as move to right across row and increases as move down column c) decreases as move to right across row and decreases as move down column d) increases as move to right across row and decreases as move down column
B
chapter 7: principle quantum number? the larger the value of n....
Characterizes the energy of the e- in a particular orbital (corresponds to bohr's energy level.) -n can be any integer greater than or = to 1 -the larger the value of n, the more energy the orbital has -energies are defined as being negative → an e- would have E=0 when it escapes the atom -the larger the value of n, the larger the orbital. -As n gets larger, the amount of energy between orbitals get smaller -determines the overall size and energy of an orbital!
chapter 7: what are complementary properties?
Complementary properties: when you try to observe the wave nature of the electron, you cannot observe its particle nature, and vice versa. Exclude one another- the more you know about one, the less you know about the other. -wave nature=interference pattern -particle nature=position, which slit it is passing through -we can never see both the interference pattern and simultaneously determine which hold the electron goes through! Can never see wave nature and particle nature at same time! Complementary properties (explained by uncertainty principle)
chapter 7: how to describe an orbital? orbitals with the same value of n are in.... orbitals with the same value of n and l are in....
Describing an orbital: each set of n, I, and mI describes one orbital. Orbitals with the same value of n are in the same principal energy level ("principal shell"). Orbitals with the same values of n and I are said to be in the same sublevel ("subshell") Summarizing sublevels: -number of sublevels in any level=n, so the n=1 level has 1 sublevel, n=2 has 2 sublevels, etc. -number of orbitals in any sublevel is=2l+1. S sublevel (l=0) has 1 orbital, p sublevel (l=1), has three orbitals, etc. -number of orbitals in a level=n^2
chapter 8: electron affinity when will you yield a larger electron affinity? electron affinity trends on periodic table?
Electron affinity of atom/ion is the energy change associated with the gaining of an electron by the atom in the gaseous state. -electron affinity defined as exothermic (-), but may actually be endothermic (+). -some alkali earth metals and all noble gases endothermic bc of repulsion -more energy that is released, larger e- affinity. -the more neg the number, the larger EA (energy going out of system) -usually negative because an atom/ion usually releases energy when it gains an electron -most groups/columns of periodic table don't exhibit any definite trend in e- affinity; among group 1A metals, however, e- affinity becomes more positive as move down column (adding e- becomes less exothermic) -electron affinity generally becomes more negative (adding e- becomes more exothermic)as move to right across period/row
Chapter 8: what are electron configurations described by? does each electron in an atom have the same set of quantum numbers, or unique sets of quantum numbers? what orbits do electrons normally occupy? How many sublevels "l" are there in comparison to n? how many orbitals "ml" are within a sublevel? how many orbitals are in a principal level? what is the maximum number of electrons in a level n?
Electron configurations described by 4 quantum numbers n, l, ml, and ms : each e- in an atom must have its unique set of quantum numbers -quantum-mechanical theory describes behavior of e-s in atoms. E-s in atoms exist in orbitals -electron configuration=description of the orbitals occupied by electrons. E- configuration for a particular atom shows the particular orbitals that e-s occupy for that atom. -ground state- lowest energy state -electrons normally occupy the lowest energy orbitals available. Number of electrons that can go into given orbital is governed by well defined rules provided by quantum mechanics: -# l sublevels within level n is n -# orbitals ml within sublevel l=2l+1 -# orbitals in principal level n=n^2 Additional concepts that affect multi-e- atoms: e- spin that is the 4th quantum #, ms=½ or ms=-1.2 and energy splitting of sublevels -maximum number of electrons in a level n = 2n^2
chapter 7: how were electrons proven to have wave nature?
Electron diffraction: beam of electrons would produce an interference pattern the same as waves do (e-s present in interference pattern). This is so even if electrons come to the slits one at a time implying that a single electron goes through both slits. -if e- beam aimed at 2 closely spaced slits, and series of detectors arranged to detect e-s after pass through slits, interference pattern similar to that of light recorded behind slits → produces interference pattern -wave nature of e- is an inherent property of individual electrons.
chapter 7: how can an electron transition to a higher energy state? what does each line in an emission spectrum correspond to?
Electron transitions: To transition to higher energy state, e- must gain correct amount of energy corresponding to difference in energy between final and initial states. -e-s in high energy states unstable, tend to lose energy and transition to lower energy states. -each line in emission spectrum corresponds to difference in energy between 2 energy states.
chapter 8: what happens when transition metals form cations? which are the first electrons to be removed? For example, Iron. When iron forms Fe2+, what does its electron configuration look like? what about Fe3+?
Electrons configurations of transition metal cations in their ground state: -when transition metals form cations, first electrons removed are valence electrons (even though other e-s added after) -e-s may also be removed from the sublevel closest to the valence shell after the valence e-s. -iron atom has two valence e-s: 1s22s22p63s23p64s23d6 -when iron forms a cation, it first loses its valence e-s: Fe2+=1s22s22p63s23p63d6 -it can then lose 3d e-s: Fe3+=1s22s22p63s23p63d5
chapter 8: what are the irregularities in electron configurations? how do you find the electron configurations for some transition metals?
Irregular electron configurations: when know that because of sublevel splitting, the 4s sublevel is lower in energy than 3d; therefore, the 4s fills before the 3d -1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f -difference in energy is not large -some of the transition metals have irregular electron configurations in which the ns only partially fills before the (n-s)d or doesn't fill at all. Therefore, their electron configuration must be found experimentally. (just know that they exist!) -irregular electron configurations: some expected and some found experimentally
chapter 7: magnetic quantum number and spin quantum number?
Magnetic quantum number, mI → orientation of orbital! -the magnetic quantum number is an integer that specifies the orientation of the orbital. The direction in space the orbital is aligned relative to the other orbitals -values are integers from -l to +l (including 0) → gives the number of orbitals of a particular shape: when I=2, the values of m1 are -2,-1,0,+1,+2 which means there are 5 orbitals with I=2 -number of orbitals for a given 2I+1 Spin quantum number (ms): specifies orientation of the spin of the e-. All e-s have same amount of spin.
chapter 8: what is metallic character? when is metallic character more prominent? If it is easy to remove an electron, is the element more or less metallic? periodic trends for metallic character? do metals or nonmetals tend to have smaller first ionization energies? what about larger electron affinities?
Metallic character- the more loosely the outer e- is bound the more metallic the character: Metallic character-how closely an element's properties match the ideal properties of a metal -easier to remove e- the more metallic the element -metallic character decreases left to right across a period (valence e-s held strongly, e- affinity increases) -metallic character increases down column (valence e-s not held as strongly) Metals generally have smaller first ionization energies and nonmetals generally have larger electron affinities (EA), except from noble gases smaller EA
chapter 8: noble gases why are noble gases especially nonreactive?
Noble gas electron configuration: noble gases have 8 valence e- (except for He which has 2). Esp nonreactive (He and Ne practically inert). The reason the noble gases so nonreactive is that electron configuration of the noble gases is esp stable → when s and p quantum sublevels completely full, overall energy of e-s that occupy that level is particularly low.
chapter 8: noble gases periodic trends for atomic radius, density, and first ionization energy? why are noble gases unreactive?
Noble gases: Atomic radius and density increase for each successive noble gas, and first ionization energy decreases -high first ionization energies of noble gases and their completely full s and p sublevels make them exceptionally unreactive
Chapter 8: what is the pauli exclusion principle? how many electrons can an orbital have? 2 eelectrons occupying the same orbital have.... maximum number of electrons in the s sublevel? p sublevel? d? f?
Pauli exclusion principle: no 2 electrons in an atom may have the same set of 4 quantum numbers: n, l, ml and ms. therefore, no orbital may have more than 2 electrons, and they must have opposite spins. -two e-s occupying the same orbital have 3 identical quantum numbers (n, l, and ml) but have different spin quantum numbers. Knowing the # of orbitals in sublevel allows us to determine the maximum # of electrons in sublevel: -s sublevel has 1 orbital: can hold 2 e-s -p sublevel has 3 orbitals: can hold 6 e-s -d sublevel has 5 orbitals; can hold 10 e-s -f sublevel has 7 orbitals; can hold 14 e-s Orbital diagrams: often represent orbital as square and electrons in orbital as arrows. Direction of arrow represents spin of electron
chapter 7: how to predict the spectrum of hydrogen? how many energy states can an e- transition to? how many lines on an emission spectrum can it generate?
Predicting the spectrum of hydrogen: -wavelengths of lines in emission spectrum of hydrogen can be predicted by calculating difference in energy between any 2 states. -e- in energy state n, there are (n-1) energy states it can transition to. Can generate (n-1) lines -bohr and quantum mechanical models can predict these lines accurately for a 1 e- system. Energy transitions in hydrogen: energy of photon released is = to the difference in energy between 2 levels the e- is jumping between -calculated by subtracting energy of initial state from energy of final state
chapter 7: what is the quantum mechanical explanation of atomic spectra? what does each wavelength in the spectrum of an atom correspond to? what happens when an atom absorbs energy? what is the final result?
Quantum mechanical explanation of atomic spectra. -each wavelength in the spectrum of an atom corresponds to an e- transition between orbitals summary: when atom absorbs energy, e- absorbs energy and is excited to unstable energy level, then light is emitted as electron falls back to lower energy level
Chapter 7: quantum mechanics and quantum mechanical model
Quantum mechanics forms the foundation of chemistry (explaining periodic table, the behavior of elements in chemical bonding, provides the practical basis for lasers, computers, and other applications) Quantum-mechanical model- explains strange behavior of electrons, and explains the manner in which electrons exist and behave in atoms. Helps us understand and predict the properties of atoms that are directly related to the behavior of the electrons (why some elements are metals and others nonmetals, why some elements gain 1 electrons when forming an ions whereas others gain 2, why some elements are very reactive and some are practically inert, why in other periodic patterns we see in the properties of the elements.)
chapter 7: radial distribution and radial distribution function? what does the radial distribution function represent? what is the total radial probability at a given radius? does the probability at a point increase or decrease with increasing distance from nucleus?
Radial distribution function: gives you an idea of where the electron is most likely to be found. Volume of shell=4pir2dr Total radial probability (at given r)=(probability/unit volume)xvolume of shell at r -radial distribution represents total probability of finding an e- within a thin spherical shell at a distance r from nucleus. -Probability at a point decreases w increasing distance from nucleus, but volume of spherical shell increases. -shape of the radial distribution function is the result of multiplying together two functions with opposite trends in r:
chapter 7: shapes of atomic orbitals? information about the s orbital?
Shapes of atomic orbitals: l=0 are s orbitals, l=1 are p orbitals, l=2 are d orbitals, and so on I quantum # primarily determines shape of the orbital (I can have integer values from 0 → n-1) -each value of I valled by particular letter that designates the shape of orbital. -s orbitals=spherical -p orbitals=two balloons tied at knots -d orbitals= 4 balloons tied at knots -f orbitals=8 balloons tied at knots I=0, the s orbital: lowest energy orbital. -each principal energy level has 1 s orbital (lowest energy orbital in principle energy state), spherical, # nodes=n-1
Chapter 8: equation for effective nuclear charge?
Shielding and penetration: -the effective nuclear charge=net positive charge that is attracting a particular electron -Z=nuclear charge, and S=number of electrons in lower energy levels. Z(effective)=Z-S is s>p>d>f Z=number protons S=number core electrons
Chapter 8: what are degenerate orbitals? what orbitals are not degenerate? what are the energies of sublevels like for multi-electron atoms? why is this? how to figure out how much energy a sublevel has? what is this energy based on?
Sublevel splitting in multi-electron atoms. -call orbitals w same energy=degenerate (ex all 3 p orbitals) -orbitals within principal level of multielectron atom not degenerate (energy depends on value of l). Energies of sublevels are split. -for multi-electron atoms, energies of sublevels are split (caused by charge interaction, and shielding and penetration) -the lower the value of the I quantum number, the less energy the sublevel has (s=(l=0) < p(l=1),d(l=2)<f(l=3)) -E(s orbital)<E(p orbital)<E(d orbital)<E(f orbital)
chapter 7: two-slit interference pattern
The diffraction of light through 2 slits separated by a distance comparable to the wavelength of light, coupled with interference, results in an interference pattern. Each slit acts as new wave source, and two new waves interfere with each other. Resulting pattern=series of bright and dark lines. An interference pattern is a characteristic of all light waves.
chapter 7: what is the electromagnetic spectrum? where is short wavelength, high frequency radiation located? where is long wavelength, low frequency radiation located? what light has the lowest energy? what has the highest energy? radio wave? gamma ray? x-rays? ultraviolet radiation? visible light? infrared radiation? microwaves?
The electromagnetic spectrum: -visible light comprises only a small fraction of all the wavelengths of light, called the electromagnetic spectrum. short wavelength, high frequency radiation on right and long-wavelength, low-frequency radiation on left. Short wavelength light has greater energy than long wavelength light: -radio wave light has the lowest energy -gamma ray light has the highest energy (and shortest wavelength) -X-rays: longer wavelengths than gamma, next to gamma. Used to image bones/internal organs -ultraviolet radiation: component of light that produces tan/sunburn. Not as energetic as gamma or x-rays, but still carries enough energy to damage biological molecules. -visible light next on spectrum: from violet (shorter wavelength, higher energy) to red. -infrared radiation: heat you feel when place your hand near hot object. -then, microwaves, then longest wavelengths radio waves.
chapter 7: what are orbitals determined by? what is the phase of an orbital? what does it determine?
The phase of an orbital: -orbitals determined from mathematical wave functions. Wave function can have + or - values (as well as nodes where wave function=0) -sign of the amplitude of a wave function=phase (look at example picture of this!) phase of wave determines how it interferes with another wave. -when orbitals interact, wave function may be in phase (same sign) or out of phase (opposite signs) → imp in bonding! look at picture to understand phases! in notes!
Chapter 7: what cannot be explained by conventional physics? what is Rutherford's nuclear model? what were the problems with his nuclear model?
The photoelectric effect as well as the observed spectra of atoms and molecules could not be explained by conventional classical physics. -atom contains a tiny dense center called nucleus. Nucleus essentially the entire mass of the atom. + charged, amount of positive charge balances the negative charge of the electrons. The electrons move around in the empty space of the atom surrounding the nucleus. Problems with nuclear model of atom: -electrons are moving charged particles (according to classical physics, moving charged particles give off energy) -therefore, e-s should constantly be giving off energy (causing atom to glow) -electrons should lose energy, crash into the nucleus, and the atom should collapse (but doesn't).
chapter 7: what is the photoelectric effect? what is this effect attributed to? -what affects the emission of electrons?
The photoelectric effect; particle nature of light: observation that many metals emit electrons when a light shines on their surface (photoelectric effect). -classical electromagnetic wave theory attributed this effect to the transfer of energy from the light to an electron in the metal, which resulted in the dislodgment of the electron. -according to this explanation, only amplitude/intensity of the light affects the emission of e-s, not the wavelength. -according to this theory, if the light is made brighter, more electrons should be ejected independent of wavelength. -the idea predicts if a dim light were used there would be a lag time before electrons were emitted (to give the electrons time to absorb enough energy)
Chapter 8: how do d block and f block elements differ?
Transition and inner transition elements: d block and f block -exhibit unique trends that differ from main-group elements -principal quantum number of d orbitals that fill across each row in the transition series=row number-1. Ex: in 4th row 3d orbitals fill, in 5th row 4d orbitals fill,etc. -as move across f block, principal quantum number of f orbitals that fill across each row=row number-2. Ex: in 6th row, 4f orbitals fill.
chapter 8: halogens periodic trends for atomic radius? ionization energy? reactivity? melting point and boiling point? density? what are the electron affinities like for halogens?
Trends in halogens: -atomic radius increases down column, ionization energy decreases down the column -very high e- affinities (good oxidizing agents to easy to reduce, very reactive so not found uncombined in nature, react with metals to form salts, compounds generally soluble in water so found in seawater.) -reactivity increases down the column. -melting point and boiling point increase down the column -density increases down the column (in general, increase in mass greater than increase in volume) -powerful oxidizing agents. -halogens all react with metals to form metal halides according to equation: 2 M + n X2 → 2 MXn M=metal, X=halogen MXn=metal halide.
chapter 7: what is visible light? what is color determined by?
Visible light- light that can be seen by human eye: Color: color of light determined by wavelength or frequency. White light is a mixture of all the colors of visible light (a spectrum... ROYGBIV). Red light=longest wavelength -when an object absorbs some of the wavelengths of white light and reflects others, it appears colored; the observed color is predominantly the colors reflected.
chapter 8: what elements tend to have irregular features? what does the number of columns in a block correspond to? how to find the number of valence electrons in an element on the periodic table? how to find the highest principal quantum number of an element using the periodic table?
d block and f block elements- D and f elements (called transition and inner transition elements) have often somewhat irregular features. -transition elements comprise d block, lanthanides and actinides (inner transition elements) comprise f block. -number of columns in a block corresponds to the maximum number of electrons that can occupy the particular sublevel of that block. Ex s block has 2 columns, to 1 s orbital can hold max of 2 e-s, p block has 6 columns (corresponding to 3 p orbitals w 2 e-s each), etc. Summary: periodic table divisible into 4 blocks corresponding to filing of 4 quantum sublevels (s,p,d,f), group number of main-group element=# valence electrons in element, row number of main group element = highest principal quantum # of that element.
chapter 7: angular momentum number
determines the shape of the orbital -angular momentum= m x v x r. -Angular momentum is conserved -l can have integer values from 0 → n-1. -each value of I is called by a particular letter that designates the shape of the orbital (s orbitals are spherical, p orbitals are like 2 balloons tied at knots, d orbitals are like 4 balloons tied at knots, f orbitals are like 8 balloons tied at the knots)
chapter 8: do ions in the same group have the same charge or different charge? does ion size increase or decrease down a group/column? how do cations compare to neutral atoms? how do anions compare to neutral atoms? a larger positive charge indicates a larger/smaller cation a larger negative charge indicates a larger/smaller anion
ions in the same group have the same charge. Ion size increases down column (higher valence shell, larger distance from nucleus). Cations smaller than neutral atoms, anions larger than neutral atoms. Cations smaller than anions (except Rb+ and Cs+ bigger or same size as F- and O2-). larger positive charge=smaller cation (for isoelectronic species (which have same electron configuration). Larger negative charge=larger anion (for isoelectronic species)
chapter 7: f orbitals?
l=3, f orbitals: each principal level with n=4 or greater contains 7 f orbitals. (ml=-3,-2,-1,0,+1,+2,+3) -4th lowest energy orbitals in principal energy state, mostly 8-lobed (some 2-lobed w toroid) -planar nodes (higher principal levels also have spherical nodes)
Chapter 7: light's wave nature and wave characteristics magnetic field vs electric field?
light has wave-particle duality. Light-a form of electromagnetic radiation. Composed of perpendicular oscillating waves, 1 for the electric field and 1 for the magnetic field. -electromagnetic radiation=a type of energy embodied in oscillating electric and magnetic fields -electric field- region where an electrically charged particle experiences a force. -magnetic field is a region where a magnetized particle experiences a force. -all electromagnetic waves move through space at the same constant speed (3.00x108m/s=speed of light=186,000 miles/s!) → how fast light waves move in a vacuum. Light travels faster than sound! -speed of sound=340 m/s=768 mph
chapter 7 what is the probability density? what is the probability distribution function? what does a high dot density near the nucleus indicate? As you move away from the nucleus, what happens to the probability density?
looks like triton^2 -Probability of finding an e- at a particular point in space for s orbital max at the nucleus (probability/unit volume) Decreases as move away from nucleus. Probability density function: represents total probability of finding an electron at a particular point in space: -in an orbital, density of dots proportional to magnitude of probability density -high dot density near nucleus indicated higher probability density for the electron -as move away from nucleus, probability density decreases -in curve comparing r to probability density, height of curve proportional to probability density (curve decreasing at decreasing rate)
chapter 7: Which of the following is not an allowed set of quantum numbers? n=4, l=3, ml=4 n=1, I=0, mi=0 n=5, l=4, ml=-2 n=2, l-1, ml=0 n=3, l=3, ml=-2
n=4, l=3, ml=4 is not allowed!
chapter 8: trends in successive ionization energies how does the removal of each successive electron change in relation to energy? how does ionization energy differ when removing an outermost electron vs when removing core electrons?
removal of each successive e- costs more energy -shrinkage in size due to having more protons than e-s -outer e-s closer to the nucleus (so harder to remove) -ionization energy increases fairly uniformly with each successive removal of an outermost electron, but then takes a large jump with the removal of the first core electron.
chapter 7: sublevels and max amount of e-s they can hold?
sublevels are s (l=0) which can hold a maximum of 2 e-s, p(l=1) which can hold a maximum of 6 e-s, d(l=2) which can hold a maximum of 10 e-s, and f(l=3) which can hold a maximum of 14 e-s.
chapter 7: a major league baseball player throws a 148.8 gram baseball at a speed of 41.4 m/s. what is the de broglie wavelength of the baseball in L? Is an electron's energy and position complementary? how can we describe an electron's position (since we can't specify its exact location at a given instant?)
use wavelength= (6.62 x 10^-34)/mv 1.08 x 10^-34 Electron energy: electron energy and position are complementary: KE=1/2mv2 -for e- states, since we can specify the energy of the e- precisely but not its location at a given instant, we can describe the electron's position in terms of an orbital (a probability distribution map showing where the e- is likely to be found)
chapter 8: when atoms form cations, what happens to the valence electrons? how much effective nuclear charge do they feel? do cations get larger or smaller as you go down a group? do they get larger or smaller as you go across a period? what about all of these questions for the formation of anions?
when atoms form cations (+ ve ions) the valence e-s are removed (more protons per e-.) These "new valence electrons" also experience a larger effective nuclear charge than the "old valence electrons", shrinking the ion even more. -traversing down a group increases the (n-1) level, causing the cations to get larger. -traversing to right across a period increases the effective nuclear charge for isoelectronic cations, causing the cations to get smaller Periodic trends in ionic radius anions: when atoms form anions, electrons are added to the valence shell (fewer protons per electron.) These "new valence electrons" experience a smaller effective nuclear charge than the "old valence electrons," increasing the size. The result is that the anion is larger than the atom. -as number of protons increase, radius for ion will get smaller. For a given number of electrons, a greater nuclear charge results in a smaller atom or ion.