Midterm 1 Flashcards
1
Q
How can you thing of electromagnetic energy
A
imagine self-propagating transverse oscillating waves of electric and magnetic fields
2
Q
how fast does electromagnetic radiation travel
A
the speed of light c = 299,792,458 ms
3
Q
what are the 3 characteristics of waves
A
- frequency, V, how many peaks pass a point per second
- Wavelength, (lambda)
- amplitude (height from centre to top
4
Q
how do you calculate the frequency or wavelength of an electromagnetic wave when you know the other
A
Frequency x wavelength = speed = 299,792,458
5
Q
what are the possible levels for n = 1
A
n = 1
l = 0
ml =0
6
Q
what are the possible levels for n = 2
A
n = 2
l = 0: nl=0
l = 1: nl = -1, nl = 0, nl = 1
7
Q
what are the levels for n = 3
A
n = 3
l = 0: ml = 0
l = 1: ml = -1, ml = 0, ml = 1
l = 2: ml = -2, ml = -1, ml = 0, ml = 1, ml = 2
8
Q
what do n, I,, and ml letters mean for orbitals
A
n = orbital level
l = suborbital (0=s, 1=p, 2=d, 3=f)
ml = different suborbital modes
9
Q
what are p orbitals like
A
- dumbbell-shaped, have an angular node passing through the nucleus
- 3 variations, x, y, z
- all p orbitals have 1 angular node all other nodes are radial
10
Q
what are d orbitals like
A
- 2 angular nodes
- 5 different orbitals within it due to ml values
- can have nodal cones (looks like a shaved ice cone)
11
Q
what is electron spin
A
electrons behave like a very small bar magnet, so the direction they spin has effects
ms = spin angular momentum quantum number
= +/- 1/2
12
Q
what are the possible spins for electrons
A
Spin up = +1/2, spin down = -1/2
these are show with up or down single sided arrows
13
Q
what is Paull’s exclusion principle
A
- no 2 electrons can have the same set of four quantum numbers
- therefor 2 electrons in the same orbital must have opposite spins
14
Q
what is the order of energy in orbitals and suborbital
A
- s<p<d<f: electron-electron repulsion causes energy of orbitals to increase within a shell in that order
- orbital size increases with n as does energy
15
Q
summarize the rules for electron orbitals
A
n = integer > 1
l = integer between 0 and n-1
ml = integer between - l and + l
ms = +/- 1/2
every orbital has n-1 nodes
16
Q
what is electron configuration
A
how electrons are distributed
17
Q
what is ground state
A
the most stable configuration or ground state is that which the electrons are in the lowest energy state
18
Q
what are the rules for writing electron configuration
A
when writing e- configuration
- fill orbitals in order of increasing energy
- no two electrons can fill the same orbital with same spin
19
Q
what is Hund’s rule
A
for degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized
20
Q
what does Hund’s rule mean
A
means to spread out electrons in the p orbital, then fill with second electrons
they fill each orbital singly with their spins parallel before any orbital gets a 2nd electron (though this can be effected by artificial means
21
Q
what is condensed electron configuration
A
electron configuration may be written in shorthand by turning the core electrons into their corresponding noble gas. (the inner shells don’t affect chemistry much)
then the valence electrons are written explicitly
(valence shell = outer shell where electrons are gained or lost in reactions)
ex:
Na 1s2 2s2 2p6 3s1 → Na [Ne]3s1
22
Q
how does the periodic table relate to orbitals and sub orbitals
A
- the period # is the value of n for the s orbital
- s-block elements = Alkali + Alkaline earth, they only have s orbitals
- p-block = group 13-18 (except He) have p orbitals being filled
- d-block group = transition metals filling in d orbital
- F-block = lanthanids + actinides, filling f block orbitals
23
Q
what are some anomalies within electron configuration
A
- Cr is [Ar]3d5 4s1 not [Ar]3d4 4s2
- Cu is [Ar]3d10 4s1 not [Ar] 3d9 4s2
this is due to the stability of half-filled and filled shell configurations
when atomic numbers are above 40 energy differences are small enough that anomalies occur
24
Q
summarize the rules for electron configuration
A
- lower energy orbitals fill with electrons first
- any orbital can hold up to 2 electrons
- if 2+ degenerate orbitals are available, one electron goes into each orbital till all are half full, then the new electrons start filling the orbits
- A particularly stable configuration is one where a set of p or d orbitals is either filled or half filled
25
what are the two methods to determine atomic size
1. half the diameter between the nuclei in a compound (bonding radius)
2. distance between stacks of molecules in a solid (non-bonding radius)
- non-bonding radius is slightly larger
- if looking at trends, pick one and stick with i
26
what are the trends in atomic size
- going down atoms get larger
- going right atoms get smaller
cesium = biggest
27
what is Z effective (as an idea)
Effective nuclear charge (Zeff) is the charge experienced by an electron on a many-electron atom
doesn’t equal the charge on the nucleus because of the effect of inner electrons
the electrons are attracted to the nucleus but repelled by other electrons which shield or screen it from the full nuclear charge
28
what is the formula for Z effective
Zeff = Z - S (screening constant or shielding constant)
Z = atomic number ie number of protons
S = number of core electrons (not valence)
29
when does Z effective go up
As you move to the right of the table the Z value goes up but the S value stays the same so Zeff goes up and the valence electrons are pulled closer to the nucleus
During the transition metals electrons are added to the core so Zeff doesn’t change
30
what happens to Zeff going down a group
Zeff gets slightly larger down a group cause screening is not perfect
31
how do d and f orbitals affect Z effective
filled d and f orbitals should be treated as core electrons
32
what is ionic radius
ionic radius is the measure used to describe the size of an ion
33
how are the sizes of cations and anions compared to atoms
- Cations are smaller than their parent, since you are taking electrons away from the outer shell, which can remove a shell. (But it also raises Zeff)
- Anions are bigger than their parent atom since the outer shell can be filled out and the total electron to electron repulsion increased.
34
when does size increase for ions of the same charge
For ions of the same charge, size increases as you go down a group
35
how does ionic radii work in isoelectric series
- members of an isoelectric series have the same number of electrons
- as nuclear charge increases the ions become smaller
- so: O2- > F- > Ne > Na+ > Mg2+ > Al3+
36
What is ionization energy
the minimum energy to remove an electron from the ground state of the isolated gaseous atom or ion
It is a positive value as the energy is required to remove electrons
37
what is I1, I2... etc
- I1 = energy to remove an electron from the gaseous atom
- I2 = energy to remove an electron from the gaseous 1+ ion
38
what are trends ionization energy
- increases across a period as Zeff going up makes it harder to remove electrons
- Decreases down a group since it is easier to remove electrons when they are further from the nucleus
- Increases for each successive electron that is removed from an element: I1 < I2 < I3
- removing a core electron takes a lot more energy than removing a valence electron
39
what are the two Main exceptions for trends in ionization energy and why do they exist
- 2 exceptions: removing the 1st p electron and 4th p electron are harder because
- 1st exception, the S electrons are more effective at penetrating shielding than p electrons
- 2nd exception, the electron to electron repulsion when they start doubling up the p orbitals makes it easier
40
what is electron configuration of ions and how do you remove the electrons order wise
- derived from the electron configuration of elements with the required number of electrons added or removed from the most accessible orbital
- Transition metals lose valence electrons s first then as many d electrons are are required to reach charge on the ion
- electrons removed from orbital with largest principle quantum number first: 4s then 3d
41
what is electron affinity
- the energy change when gaseous atoms gain an electron to form an ion. Opposite to ionization energy
- Ea is negative for stable anions since energy is usually released when an atom gains an electron
42
what are trends in electron affinity
- don’t change much as you move down a group. Since doing down the attraction of the electron to the nucleus is less, but so is the electron electron repulsion, so it balances out
- becomes more exothermic as we move left to right across a row as Zeff goes up
- there are discontinuities when entering a new subshell, or pairing electrons in p orbitals
43
what are chemical bonds
- When atoms and ions are strongly attracted to each other they can bond
- in chemical bonds, electrons are transferred or shared between atoms
Valence electrons are involved in bonding
44
what are the three types of chemical bonding
- In Ionic bonding the electrons are transferred, then electrostatic attraction leads to compounds like table salt
- in Metallic bonds, electrons delocalize throughout a lattice of closely packed metal ions and are shared over a long distance
- In covalent bonds electrons are shared between different atoms
45
what do lewis symbols do
show an elements symbol surrounded by dots representing the valence electrons
46
what is the octet rule
- an octet is full of s and p subshells
- takes 8 electrons to fill up
- a full octet is what noble gases tend to have, and what atoms tend to approach during bonding
the octet rule states that atoms tend to gain, lose, or share electrons till they are surrounded by 8
47
how does Zeff affect bonding
Since Zeff is highest on the right side of the table, and ionization energy is lowest on the left. Atoms from the left side usually lose electrons and those on the right side gain them
48
how does ionic bonding work
- metal reacts with non metal to fill octet
- the metal losses electrons and becomes a cation
- the non metal gains electrons to become an anion
49
what structure do ionic substances have
- ionic compounds have lattice structure
- ex in NaCl: each Na+ cation is surrounded by 6 Cl- anions and vice versa
50
how do covalent bonds work
- atoms share electrons, several electrostatic interactions in these bonds
- Nuclei and electrons attracted to each other
- Nuclei and nuclei repel
- electrons and electrons repel
51
what happens when atoms approach to bond covalently
As 2 atoms approach to form a covalent bond there is a point (the equilibrium bond length) where the attractive forces between opposite charges balance the repulsive forces from the like charges
52
what are polar and non-polar covalent bonds
- in nonpolar bonds like F2 electrons are shared equally
- in polar bonds the electrons are unequally shared (these are more common)
- One atoms attracts the bonding electrons more than the other
53
What is Electronegativity
the ability of an atom in a molecule to attract electrons to itself
related to ionization energy and electron affinity
54
how does electronegativity vary
higher electronegativity to the right and up
from 0.7 (Cs) to 4.0 (F)
55
how does electronegativity affect bonding
- the bigger the difference in electronegativity is between the 2 atoms the more polar the bond
- If the difference is large enough an ionic bond will form cause the electron is stole
- if above 2 than ionic, if 0.5-2.0 than polar covalent, if 0-0.5 than non-polar covalent.
56
what is the dipole moment
- polar molecules have a centre of + charge and a centre of - charge that don’t coincide.
- the dipole moment (µ) is produced by opposite charges separated by a distance (r)
- µ = Qr (in debyes (D))
57
how do lewis structures work for compounds
- show the valence electrons, bonding or non bonding
- (—) shows bonding (**) shows non bonding pair
- multiple bonds are shown by multiple lines. (double bond = two shared pairs) (triple bond = 3 shared pairs)
58
59
how do you make the lewis structure for a compound
1. find the sum of valence electrons of all atoms in the polyatomic ion of molecule
If its an anion, add one electron for each negative charge. Opposite for cations
2. The central atom is usually the less electronegative one
draw a line from it to the outer electrons
3. fill the octets of the outer atoms with non bonded pairs
4. fill the octet of the central atom with non bonded pairs
5. If you run out of electrons before the central atom has an octet, change some of the non bonded pairs into double or triple bonds.
6. Make all possible Lewis structure that do this
60
how do you choose the best lewis structure for a molecule
- assign formal charges to each atom by counting the number of electrons in lone pairs and half the electrons in bonds it has. The subtract the number of valence electrons for that atom.
- Its best to have the fewest absolute number of charges
- Its best to have the negative charge on the more electronegative atoms
- choose whichever lewis structure best fits these two criteria
61
what are the exceptions to the octet rule
1. those with an odd number of electrons
relatively rare and unstable. All atoms except one have an octet
2. those with less than an octet
Sometimes using a double bond to fill the octet of the central atom will put a positive charge on the more electronegative atom (like in BF3). Don’t fill the octet of the central atom in these cases.
Usually boron or aluminum
3. those with more than 8 valence electrons
atoms like PCl5 can only exist is P has 10 electrons around it.
Atoms from the third period and beyond can accomodate more than an octet
62
what is resonance structure and how do you draw it
for some molecules like NO3- the lewis structure has a double bond, but the observed structure of nitrate has all three N-O bonds being the same length and all Os having the same partial negative charge
for Nitrate it resonates between three possible structures (each with a different O-N having the double bond)
it is drawn with each O-N having a full bond, and a dotted line showing a partial bond, and a -0.67 charge on each oxygen to represent the average structure between the three.
63
What is Bond enthalpy
determined by how much energy is required to break the bond: the bond enthalpy (D)
positive since bond breaking is endothermic
- multiple bonds take more energy to break
- most D values are averages due to differences in different molecules
64
how stable are ionic bonds
ionic compounds are stable because of the electrostatic attraction between positive and negative ions
65
how is the change in enthalpy found for a reaction
the ∆H for a reaction is determined by comparing the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed
66
what is the idea of lattice energy
- Lattice energy is the energy required to completely separate.a mole of a solid ionic compound into its gaseous ions
- Coulomb’s law governs the energy associated with electrostatic interactions
67
what is the formula for lattice energy
- Eel = KQ1Q2/d
- K = 8.99 x 10^9 J m/C^2
- Q1, Q2 = charges. d = distance
- the values are positive as it the energy required to break the lattice of the ionic solid.
68
what energies all play a factor in a reaction
Ionization energy, electron affinity, and lattice energy give us a good idea of the energies involved in the formation of an ionic compound from its elements. Along with the energy required to get the elements from their normal state to gas-phase atoms
A born-Haber cycle is used to analyze factors contributing to stability of ionic compounds.
69
what is molecular shape
while Lewis structures show how atoms are connected, too describe a compounds shape we use its bond angles and lengths (in picometers).
A molecules shape plays an important role in its reactivity
70
what is VESPER
VESPER = Valence Shell Electron Pair Repulsion
by assuming valence electrons are placed as far as possible form each other we can predict molecular shapes
VESPER says that electron domains should be as far apart as possible to minimize repulsion
71
what are electron domains
Electron domain = region occupied by electrons (nonbonding, single bonds, multiple bonds all equal an electron domain)
72
what are the basic shapes of molecular geometry for 2-6 atoms
Basic shapes for molecules with no non-bonded pairs. Also electron domain geometry for number of electron domains
2 = Linear: 180º between the bonds
3= Trigonal planar: 120º between the bonds
4= Tetrahedral: 109.5º between bonds
5= Trigonal bipyramidal: 90º or 120º between bonds
6= Octahedral: 90º between bonds
73
how do you find the electron domain geometry
1. draw lewis structure, count domains
2. arrange electron domains to minimize repulsion
3. inspect arrangement of atoms to determine geometry
74
are non bonding pairs bigger than bonding pair, or opposite
non bonding pairs are physically larger than bonding pairs so they tend to decrease bond angles in a molecule (between bonds)
75
what are the possible molecular geometries for a trigonal planar electron geometry
0 non bonding pairs = trigonal planar
1 non bonding pair = bent
76
what are the possible molecular geometries for a tetrahedral electron geometry
0 non bonds = tetrahedral
1 non bond = trigonal pyrmaidal
2 non bonds = bent
77
what are the possible molecular geometries for a trigonal bipyramidal electron geometry
0 non bonds = trigonal bipyramidal
1 non bond = see-saw
2 non bonds = T shaped
3 non bonds = linear
78
what are the possible molecular geometries for a Octahedral electron geometry
0 non bonds = octahedral
1 non bond = square pyramidal
2 non bonds = square planar
rest not known
79
how are the lone pairs positioned in square planar
opposite sides of the central atom
80
where do the non bonds go on a molecule with trigonal bypyramidal electron geometry
on the equatorial electron domains
81
what are the two types of electron domains on a trigonal bipyramidal molecule
axial and equatorial
82
how do you determine the shapes of larger molecules
take each section one at a time and do it normally
83
What kind of molecules interact with electric fields
those that are polar
84
whan is a binary compound polar
when the negative and positive charge do not coincide
85
are the 5 parent ABn molecules polar
no, all are non-polar. though they can be unbalanced due to differentiators or lone pairs
86
what is the order of electromagnetic radiation in order of shortest to longest wavelength
gamma, xray, UV, visible light, Infrared, terahertz, microwave, radio
87
what term describes how waves interact
interference
88
what are the two types of interference
constructive - add to each other
destructive - subtract from each other
can also be more complicated
89
what are standing waves
waves confined within some region of space
90
what happens if both ends of a standing wave are fixed
then only a half integer number of wavelengths can form
91
what do more nodes in a wave lead to
more energy
92
what do hot objects emit
light, with wavelength dependent on temp
93
what did Max Planck propose
a theory explaining that energy can only be absorbed or released in certain amounts called quantums
94
what is the photoelectric effect
an effect where a metal surface absorbs light and emits electrons for certain frequencies of shorter wavelength light
shorter wavelength = faster electron
95
what did Einstein propose
that light could have partial like properties which he called photons (thought of as wave packets of light)
he also said that metal surfaces absorb 1 quantum of energy from the light, and the remaining is turned to KE and used to shot off an electron
96
what are line spectra
monochromatic is 1 wavelength of light, but in continuous radiation multiple wavelengths make a continuous spectrum.
Excited gaseous elements produce a line spectrum of different light wavelengths
97
what were Bohrs 3 postulates
1. only orbits of specific radii are permitted for electrons in atoms
2. an electron in an orbit has a specific energy
3. energy is emitted or absorbed as an electron moves to different energy states
98
when does photoemission occur
when an electron drops to a lower level
99
how do you calculate the energy of a level
En = -(HCR) (1/n^2) Joules
h = plancks constant
c = speed of light
R = ideal gas constant
n = level
turns to :
En = -2.18x 10^-18 (1/n^2) joules
100
how do you calculate the change in energy from electrons changing levels
∆E = -2.18x10^-18 (1/nf^2 - 1/ni^2) joules
101
who discovered the electron
Sir Joseph John Thomson
102
who showed that the electron acts as a standing wave
Sir George Paget Thomson (JJT Thomson's son)
103
what are matter waves
the wave characteristics of material particles.
Suggested by Louis de Proglie
Matter has such small wave lengths it does nothing
104
what can electrons be thought of as
standing waves, and since only they must fit the right number of nodes, only certain wavelengths can exist
105
What does the dual nature of matter set a limit on
on how precisely we can know the location and momentum of an object. measuring impacts what is being measured
the more precisely we know the speed, the less we know the position
106
what did schrodinger propose
an equation containing both wave and particle terms, its solution is known as the wave function and describes the behaviour of a quantum object
the wave function squared = probability density of where the particle is
107
what does the probability density give us for electrons around an atom
3 quantum numbers
108
what is n
the principle quantum number
must be a positive integer 1, 2, 3, 4
tells which electron shell the electron is in
109
what is I
the angular momentum number
describes the shape, specifies subshell for electron
max = n-1
0 = s, 1 = p, 2 = d, 3 = f
110
what is Mi
the magnetic quantum number
designates specific orbital and orientation
can be -I to +I
111
