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Test 4 study Flashcards

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1
Q

What energy does the electron of an H atom have?

A

E = -2.178 x 10^-18 J (1/n^2)
- E = energy
- n = energy level or principal quantum number (1, 2, 3, …)

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2
Q

How can electron transitions in H atoms be calculated?

A

Delta E = energy of final state - energy of initial state
- E = -2.178 x 10^-18 J (1/n^2 f) - [-2.178 x 10^-18 J
(1/n^2 f)]
- E = -2.178 x 10^-18 J (1/n^2 f - 1/n^2 i)

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3
Q

What is the Bohr model?

A
  • That an electron can only have specific amounts
    of energy
  • Electrons travel in orbits at fixed distances from
    the nucleus
  • The electron emits radiation moving from a
    higher to a lower energy orbit. The distance
    between the orbits determines the energy of the
    radiation
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4
Q

What is a wavefunction and what are the three main parts of it?

A

Wavefunction - gives the probability of finding an electron in a volume of space, used to generate orbitals, has three quantum numbers (variables)
1. Principal quantum number, n - orbital energy,
size
2. Angular momentum quantum number, l -
orbital shape
3. Magnetic quantum number, ml - orbital
orientation

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5
Q

What is principal quantum number n? What does it correspond to?

A
  • Characterizes the energy of an electron in a
    given orbital
  • Larger values of n correspond to higher orbital
    energy and larger atomic size
  • n>/= 1; n=0 is the nucleus
  • The largest n of any ground state atom is 7
  • As n gets larger, the difference between energy
    levels gets smaller
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6
Q

What is angular momentum quantum number l?

A
  • Angular momentum quantum numbers (l) have
    numeric and letter values
    - 0 s, 1 p, 2 d, 3 f, 4 g
  • A set of orbitals with the same n value is a shell
    or level
  • A set of orbitals with the same nl value is a
    subshell/sublevel
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7
Q

What are forbidden orbitals?

A
  • For any orbital, n>1 must always be true
    - n = 1, 2, 3, …
    - l = 0, …, (n-1)
  • no n </= l combinations exist
    • 1p, 1d, 2d, 1f, 2f, 3f, 1g, 2g, 3g, 4g
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8
Q

What is magnetic quantum number ml?

A
  • Distinguishes the orbitals available within a
    subshell
  • The set of ml values for any subshell is -1, …, 0,
    …, 1
  • Subshells always have to have an odd number of
    orbitals
    - l=0 has 1, l=1 as 3, l=2 has 5, l=3 has 7, l=4
    has 9
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9
Q

What does the Uncertainty Principle do?

A

The Uncertainty Principle limits what we know about electron movement
- The best we can do is describe a region with a
high probability of finding an electron using
wavefunctions
- Solving wavefunction gives a 3-D scatterplot,
90% of the points make up an orbital

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10
Q

What is an electron density map?

A

Electron density maps are like time-elapsed photos of an electron’s position
- They are darker where the electron is more
often found
- The probability of finding an electron at the
nucleus is zero

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11
Q

Is it accurate to think of orbitals as hard containers that trap electrons?

A

No, don’t think of orbitals as hard containers that trap the electron
- 10% of the time, the electron is outside them
- There’s a 90% probability an electron will be in
the orbital

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12
Q

s Orbitals

A
  • s orbitals are spherical and centered at the
    nucleus (l=0)
  • as n gets bigger, orbital size and energy increase
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13
Q

What are nodes?

A

Nodes - planar or spherical surfaces where wavefunctions change sine
- Except for 1s, wavefunctions change sign (+/-) at
least once
- The signs have nothing to do with charge
- Electrons are never found at a node
- For any orbital, numbers of nodes = (n-1)

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14
Q

What are radial distribution functions?

A

These represent the probability of finding an electron at a given distance from the nucleus
- This probability drops off at and far from the
nucleus
- At large n’s electrons are found further from the
nucleus

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15
Q

p Orbitals

A
  • n >/= 2 shells have three degenerate p orbitals
    • ml = -1, 0, +1
  • Each is centered along a different axis
  • They are bilobal (dumbbell-shaped) with a node
    at the nucleus (l=1)
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16
Q

What are degenerate orbitals?

A
  • Hydrogen: orbitals with the same value of n are
    degenerate (equal in energy(
  • Other atoms: oly orbitals with the same n and l
    are degenerate
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17
Q

d Orbitals

A

n >/= 3 , shells have five degenerate d orbitals (l=2)
- ml = -2, -1, 0, +1, +2

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18
Q

f Orbitals

A

n >/= 4, shells have seven degenerate f orbitals (l=3)
- ml = -3, -2, -1, 0, +1, +2, +3

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19
Q

g Orbitals

A
  • g orbitals exist in n >/= 5 shells
  • No element uses them in the ground state; the
    lightest would be Z = 121
  • Has nine subshells, ml = -4, -3, -2, -1, 0, 1, 2, 3, 4
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20
Q

What is an easy way to identify orbitals?

A
  • Dumbbell shape = p, 1 node, n=2, 2p
  • Cloverleaf shape = d, 2 nodes, n=3, 3d
  • Sphere shape = s, 3 nodes, n=4, 4s
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21
Q

What is a ground state?

A

All electrons are in lowest energy orbitals possible

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22
Q

What is electron configuration?

A

A notation for distribution of electrons into the orbitals of a ground state atom
- The number is n, the leter is l, the superscript is
the number of electrons in that sublevel
- Eg. Li: 1s^2 2s^1

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23
Q

What are the two ways to write electrons configurations?

A

Examples using configuration of vanadium
- Regular: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^3
- Empirical gas: [Ar] 4s^2 3d^3

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24
Q

What is an orbital diagram?

A

Electrons are drawn as arrows, orbitals as boxes
- Same data as electron configuration plus
electron spin
- An orbital can hld 0, 1, or 2 electrons
- If it holds 2, the spins must be opposed

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24
What is an orbital diagram?
Electrons are drawn as arrows, orbitals as boxes - Same data as electron configuration plus electron spin - An orbital can hold 0, 1, or 2 electrons - If it holds 2, the spins must be opposed
25
What is an orbital diagram?
Electrons are drawn as arrows, orbitals as boxes - Same data as electron configuration plus electron spin - An orbital can hld 0, 1, or 2 electrons - If it holds 2, the spins must be opposed
26
What is Hund's rule?
- Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied - Electrons in singly occupied orbitals have parallel spins - It is not forbidden to have two electrons occupying orbitals while there are still blank orbitals but it means the electrons are in an excited state
27
What group of elements contains electron configuration exceptions?
- Main group elements don't have exceptions to their electron configurations - Transition metal electron configurations are rife with exceptions
28
What is the electron spin quantum number ms?
- This quantum number has two values: +/- 1/2 (or up and down arrow) - The values refer to two spin vectors of the electron
29
What is the Pauli exclusion principle?
- In an atom, no electrons can have the same four quantum numbers (n, l, ml, ms) - An orbital holds at most two electrons; if it has two, their spins are opposed
30
What are the four groups of quantum numbers?
1. n - specifies a shell 2. n, l - specifies a subshell 3. n, l, ml - specifies an orbital 4. n, l, ml, ms - specifies an electron
31
What are the dimensions of the periodic table in relation to orbitals?
The number of orbitals in a sublevel determines the maximum number of electrons it can hold - s sublevels: 1 orbital, 2 electrons max - p sublevels: 3 orbitals, 6 electrons max - d sublevels: 5 orbitals, 10 electrons max - f sublevels: 7 orbitals, 14 electrons max
32
What are multielectron atoms?
- Hydrogen: the sublevels in each principal energy level are degenerate: ns=np=nd=nf - Other atoms: sublevel energies are split due to electron-electron repulsion, a sublevel with a lower l value has lower energy: ns
33
Why is a 1s orbital lower in energy than a 7s?
- 1s electrons are closer to the nucleus - 1s electrons experience no shielding, 7s electrons have 85 lower energy electrons between them and the nucleus
34
In which orbital do electrons have a greater probability of being closer to the nucleus, a 2s or 2p electron?
2p, even though it's higher in energy
35
How do you fill a subshell?
- List each energy shell on a row, writing the subshells (s, p, d, f) by increasing energy - For the order the subshells are filled, draw parallel upper right to lower left diagonals - Or, follow the atomic numbers on the periodic table
36
What are valence electrons?
Bonding electrons - Main-group elements: valence electrons in orbitals with the largest value of n - Transition metals: ignore - Numbering the eight main group families left to right gives the number of valence electrons in that group - As has 5 valence electrons
37
What are core electrons?
- Not valence electrons - Not used in bonding - Noble gas electrons are core electrons, hence the abbreviation
38
By looking at two atoms on the periodic table, you can predict which has the higher what?
1. Atomic radius 2. Ionic radius 3. Ionization energy 4. Electron affinity 5. Metallic character 6. Many others
39
What is an atomic radius?
Half the distance between nuclei in a molecule of identical atoms - Atomic radius increases down a group - Larger n = larger size - Atomic radius decreases across a period (left to right) - Moving right across a period, nuclear charge increases but shielding doesn't, Z eff increases as electrons move in
40
What is effective nuclear charge (Z eff)?
The nuclear charge an electron fels considering shielding effects - Core electrons strongly shield valence electrons - Valence electrons poorly shield each other - Electrons are attracted to the nucleus but repelled by each other
41
What is the usual configuration for elements that only form one ion?
For elements that only form one ion, that ion usually has a noble gas configuration
42
What is the electron configuration of metals?
- All metals first lose electrons from the sublevel with the highest n value - For transition metals electrons are lost and filled in different orders
43
What are the magnetic properties of atoms and ions?
- Paramagnetic species have >/= 1 unpaired electron(s) and attracted to magnetic fields - Diamagnetic species have no unpaired electrons and are slightly repelled by magnetic fields - Easy to tell by consulting orbital diagrams - Eg. silver is paramagnetic
44
What are the trends in ionic radius?
- Ionic radii increase going down a group - Cations are smaller than its neutral atom; anions are bigger - Na+ < Na - Cl < Cl- - In an isoelectronic (same number of electrons) series: - Larger positive charge = smaller cation - Larger negative charge = larger anion
45
How do cations form?
- Cations are smaller than their neutral atoms - Cations form by electron loss; the decrease in electron repulsions pulls in the remaining electrons
46
How do anions form?
- Anions are larger than their neutral atoms - Anions form by electron gain; an increase in electron repulsions causes electrons to spread out
47
What is ionization energy?
- The energy required to remove an electron from an atom or ion - M(g) --> M+ (g) + e- - Always endothermic - Measures how strongly an atom holds its electrons
48
What are the trends in ionization energy?
- IE increases up a group because valence electrons are closer to the nucleus - It takes less energy to remove an electron farther from the nucleus - IE generally increases to the right as Z eff increases - The larger an electron's Z eff, the more energy it takes to remove it
49
What are some characteristics of alkali metals?
- Alkali metals are only shiny when freshly cut or stored under mineral oil or argon - Alkali metals (Li, Na, K, Rb, Cs) have very low IEs - To react them with water is not an IE, but either way they lose an electron - M(g) --> M+ (g) + e- (IE) - M(s) --> M+ (aq) + e- (H2O rxn)
50
What are the irregularities in the IE trend?
- Ionization energy generally increases from left to right across a period with a few exceptions - Which is easier to remove: B or Be?
51
What are the irregularities in the IE trend?
- Ionization energy generally increases from left to right across a period with a few exceptions - Which is easier to remove: electron from B or Be? - Removing from Be disrupts a full sublevel - Removing from B creates a full sublevel - Which is easier to remove: electron from N or O? - Removing from N disrupts a half-full sublevel - Removing from O creates a half-full sublevel
52
What is successive ionization energies?
- Once ionized more energy will remove a second electron, then a third, and so on - IE1 < IE2 < IE3 - Outer electrons get closer to the nucleus and are harder to remove - There is a smaller increase in energy for each successive valence electron removal and a large increase for core electrons
53
What is electron affinity?
- Electron affinity (EA) - the energy change when an atom gains an electron: M(g) + e- --> M-(g) - Usually exothermic, sometimes endothermic - There is no strong EA trend, in any period: - Noble gases have the most endothermic EA - Halogens have the most exothermic EA
54
What are the three main types of chemical bonds?
- Covalent bond - sharing of electrons between two non-metal nuclei - Ionic bond - transfer of electrons between a metal and a nonmetal atom - Metallic bonding - why most metals are solid
55
What is a lewis dot structure?
- LDS represent valence electrons as dots around the atom - H follows the duet rule, it forms stable compounds by sharing 2 electrons - Most main-group atoms follow the octet rule: they tend to have 8 electrons (bonds + lone pairs) in their valence shell
56
What two types of bonds quantify bond strength?
- Ionic bonds - lattice energy - Covalent bonds - bond energies
57
What is lattice energy?
- Lattice energy- the delta H when gas phase ions form one mole of an ionic solid: - M+ (g) + X- (g) --> 1MX (s) - Always exothermic - More exothermic LE = stronger ionic bond - Lattice energy isn't measured in a lab but calculated using Hess' Law - The sum of all the reactions equals the standard enthalpy of formation of the ionic compounds - One of the summed reactions is lattice energy
58
What are polar covalent bonds?
- For some covalent bonds, electrons are not equally shared, such polar covalent bonds can be shown as: - Partial charges - Dipole; arrow points to more electronegative atom - Electrostatic potential map electron-rich and electron-poor regions
59
What is electronegativity?
The ability of an atom to attract shared electrons to itself in a chemical bond - Delta EN can be used to predict bond type: - Small EN = covalent, medium EN = polar covalent, large EN = ionic
60
What are some common characteristics of ionic compound?
- Large delta EN - Cation-anion arrays - Solids at STP - High melting/boiling points - Conduct as liquids
61
What are some common characteristics of covalent compound?
- Small delta EN - Discrete molecules - Solid/liquid/gas at STP - Low melting/boiling points - Insulating as liquids

Chemistry (9 decks)

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