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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
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 hold 0, 1, or 2 electrons
- If it holds 2, the spins must be opposed

25
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

26
Q

What is Hund’s rule?

A
  • 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
Q

What group of elements contains electron configuration exceptions?

A
  • Main group elements don’t have exceptions to
    their electron configurations
  • Transition metal electron configurations are rife
    with exceptions
28
Q

What is the electron spin quantum number ms?

A
  • This quantum number has two values: +/- 1/2 (or
    up and down arrow)
  • The values refer to two spin vectors of the
    electron
29
Q

What is the Pauli exclusion principle?

A
  • 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
Q

What are the four groups of quantum numbers?

A
  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
Q

What are the dimensions of the periodic table in relation to orbitals?

A

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
Q

What are multielectron atoms?

A
  • 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<np<nd<nf
    - As z increases: electron-nucleus attraction
    increases, the number of electrons increases
    electron-electron repulsions increase, outer
    electrons are “shielded” by inner electrons
33
Q

Why is a 1s orbital lower in energy than a 7s?

A
  • 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
Q

In which orbital do electrons have a greater probability of being closer to the nucleus, a 2s or 2p electron?

A

2p, even though it’s higher in energy

35
Q

How do you fill a subshell?

A
  • 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
Q

What are valence electrons?

A

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
Q

What are core electrons?

A
  • Not valence electrons
  • Not used in bonding
  • Noble gas electrons are core electrons, hence
    the abbreviation
38
Q

By looking at two atoms on the periodic table, you can predict which has the higher what?

A
  1. Atomic radius
  2. Ionic radius
  3. Ionization energy
  4. Electron affinity
  5. Metallic character
  6. Many others
39
Q

What is an atomic radius?

A

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
Q

What is effective nuclear charge (Z eff)?

A

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
Q

What is the usual configuration for elements that only form one ion?

A

For elements that only form one ion, that ion usually has a noble gas configuration

42
Q

What is the electron configuration of metals?

A
  • 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
Q

What are the magnetic properties of atoms and ions?

A
  • 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
Q

What are the trends in ionic radius?

A
  • 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
Q

How do cations form?

A
  • Cations are smaller than their neutral atoms
  • Cations form by electron loss; the decrease in
    electron repulsions pulls in the remaining
    electrons
46
Q

How do anions form?

A
  • Anions are larger than their neutral atoms
  • Anions form by electron gain; an increase in
    electron repulsions causes electrons to spread
    out
47
Q

What is ionization energy?

A
  • 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
Q

What are the trends in ionization energy?

A
  • 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
Q

What are some characteristics of alkali metals?

A
  • 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
Q

What are the irregularities in the IE trend?

A
  • Ionization energy generally increases from left to
    right across a period with a few exceptions
  • Which is easier to remove: B or Be?
51
Q

What are the irregularities in the IE trend?

A
  • 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
Q

What is successive ionization energies?

A
  • 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
Q

What is electron affinity?

A
  • 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
Q

What are the three main types of chemical bonds?

A
  • 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
Q

What is a lewis dot structure?

A
  • 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
Q

What two types of bonds quantify bond strength?

A
  • Ionic bonds - lattice energy
  • Covalent bonds - bond energies
57
Q

What is lattice energy?

A
  • 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
Q

What are polar covalent bonds?

A
  • 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
Q

What is electronegativity?

A

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
Q

What are some common characteristics of ionic compound?

A
  • Large delta EN
  • Cation-anion arrays
  • Solids at STP
  • High melting/boiling points
  • Conduct as liquids
61
Q

What are some common characteristics of covalent compound?

A
  • Small delta EN
  • Discrete molecules
  • Solid/liquid/gas at STP
  • Low melting/boiling points
  • Insulating as liquids

Chemistry (9 decks)

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