Test 4 study Flashcards
What energy does the electron of an H atom have?
E = -2.178 x 10^-18 J (1/n^2)
- E = energy
- n = energy level or principal quantum number (1, 2, 3, …)
How can electron transitions in H atoms be calculated?
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)
What is the Bohr model?
- 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
What is a wavefunction and what are the three main parts of it?
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
What is principal quantum number n? What does it correspond to?
- 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
What is angular momentum quantum number l?
- 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
What are forbidden orbitals?
- 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
What is magnetic quantum number ml?
- 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
What does the Uncertainty Principle do?
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
What is an electron density map?
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
Is it accurate to think of orbitals as hard containers that trap electrons?
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
s Orbitals
- s orbitals are spherical and centered at the
nucleus (l=0) - as n gets bigger, orbital size and energy increase
What are nodes?
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)
What are radial distribution functions?
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
p Orbitals
- 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)
What are degenerate orbitals?
- 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
d Orbitals
n >/= 3 , shells have five degenerate d orbitals (l=2)
- ml = -2, -1, 0, +1, +2
f Orbitals
n >/= 4, shells have seven degenerate f orbitals (l=3)
- ml = -3, -2, -1, 0, +1, +2, +3
g Orbitals
- 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
What is an easy way to identify orbitals?
- Dumbbell shape = p, 1 node, n=2, 2p
- Cloverleaf shape = d, 2 nodes, n=3, 3d
- Sphere shape = s, 3 nodes, n=4, 4s
What is a ground state?
All electrons are in lowest energy orbitals possible
What is electron configuration?
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
What are the two ways to write electrons configurations?
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
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
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
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
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
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
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
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
What are the four groups of quantum numbers?
- n - specifies a shell
- n, l - specifies a subshell
- n, l, ml - specifies an orbital
- n, l, ml, ms - specifies an electron
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
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<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
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
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
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
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
What are core electrons?
- Not valence electrons
- Not used in bonding
- Noble gas electrons are core electrons, hence
the abbreviation
By looking at two atoms on the periodic table, you can predict which has the higher what?
- Atomic radius
- Ionic radius
- Ionization energy
- Electron affinity
- Metallic character
- Many others
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
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
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
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
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
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
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
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
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
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
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)
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?
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
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
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
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
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
What two types of bonds quantify bond strength?
- Ionic bonds - lattice energy
- Covalent bonds - bond energies
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
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
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
What are some common characteristics of ionic compound?
- Large delta EN
- Cation-anion arrays
- Solids at STP
- High melting/boiling points
- Conduct as liquids
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
