Unit 1 Flashcards
1
Q
frequency
A
peaks per second in hertz (Hz) or s^-1… variable is v
2
Q
speed of light
A
assuming a vacuum for this course… variable is c
3
Q
wavelength
A
length from peak/trough of wave. can be nanometers or meters for units (use 10^9/10^-9 as conversion)
4
Q
Electromagnetic Radiation Spectrum
A
shorter wavelength, higher frequency, higher energy
5
Q
c= v(lambda)
A
inversely proportional c= speed of light constant v= frequency (waves per second) lambda= wavelength - don't forget to convert wavelength to match units of light
6
Q
energy levels
A
are called shells, orbits or symbolized by “n”. lower orbitals have less energy ex (n= 1 has less energy than n=4)
7
Q
distance between energy levels
A
- energy levels get closer together the further you go
8
Q
electron goes down an orbital
A
- electron loses energy so it will emit light of wavelength equal to energy lost
9
Q
electron goes up an orbital
A
electron absorbs energy equal to wavelength of light given to it
10
Q
absorption
A
electron goes up shell due to energy given to it
11
Q
emission
A
electron goes down shell due to energy emitted
12
Q
continuous spectrum of light (line spectrum)
A
shows rainbow
- all wavelengths on spectrum are visible light
13
Q
emission lines (line spectrum)
A
few colours on a black backdrop visually
- the colours are wavelengths of light emitted when a gas atom was originally excited but that then electron emits the light was it goes back to rest state
14
Q
absorption (line spectrum)
A
while light is going through a gas sample of an element, the electron would absorb some energy of that light causing
15
Q
at what speed does all electromagnetic radiation move at?
A
the speed of light
16
Q
electromagnetic radiation and energy transfer
A
ex. when molecules absorb radiation, it increases the energy of the molecules causing more collisions and a rise in temperature (case of microwave at least)
17
Q
planck’s quantum theory
A
energy can only be gained or lost in whole number multiples of hv where h is a constant. can find energy absorbed or released
18
Q
quantum
A
small packet of energy that can only occur in discrete units. system can only transfer energy in whole quanta thus energy seems to have particulate properties.
19
Q
photosns
A
einstein discovered that electromagnetic radiation can be seen as a stream of particles called photons
20
Q
delta e=nhv OR delta e=hc/(lambda)
A
gives the amount of a single quantum (the energy of a photon of light)
n= an integer h= planck's constant v= frequency
21
Q
dual nature of light
A
electromagnetic radiation can show certain characteristics of particulate matter which is called dual nature
22
Q
de Broglie’s equation
A
lambda= h/mv
23
Q
lambda= h/mv
A
used to calculate the wavelength for a particle
- h=planck’s constant
m= mass
v= AHAH it is velocity
24
Q
1 joule equals
A
1 kg times m^2 / s^2
25
hydrogen emission line spectrum
shows that only certain energies are allowed for the electron in the hydrogen atom AKA the energy of the electron in hydrogen is quantized. changes in energy between discrete energy levels in hydrogen will produce only certain wavelengths of emitted light. if any energy level was allowed, the emission spectrum would be continuous
26
emission spectrum
emission: when hydrogen molecules absorb energy and some of the bonds are broken. this results in excitation of electrons which means they contain excess energy which they release by emitting light of various wavelengths to produce the emission spectrum
contains black but with some colour bars in between
27
if energy levels in atoms were not quantized
all light would be white (contains all wavelengths)
28
delta e= -2.178 times 10^-18 J times z^2 (1/n^2f-1/n^2i)
```
n= integer. larger n is larger orbit radius
z= atomic number (usually 1 for hydrogen)
f= final
i= initial
```
this is negative when there is emission
29
e= -2.178 times 10^-18 J times (z^2/n^2)
gives the energy of each energy level
30
n=6 to n=1
n=1 has more negative energy because the electron in that level is more tightly bound to the smallest allowed orbit. the change in energy then is negative as energy is lost and electron is now in a more stable state. the energy is carried away from the atom by the production (emission) of a photon
31
bohr model
1. model fits the quantized energy levels of the hydrogen atom and postulates only certain allowed orbits for the electron
2. as electron becomes more tightly bound, its energy becomes more negative relative to the zero-energy reference state.
32
quantum model vs bohr
Bohr
Shells have set radii and electrons orbit the nucleus on those shell surfaces
Each energy level is a single shell
The shells surround the nucleus like planets surround the sun
Based on one number, the energy level, n
Quantum model
Each energy level is a collection of different orbitals/subshells (whose shape is defined by where you are most likely to find the electron 90% of the time)
Each orbital can hold 2 electrons
The orbitals have different complex shapes and electrons can travel anywhere within the space contained by the orbital
Each electron in an atom can be uniquely identified by 4 quantum numbers
33
quantum numbers
N
L
Ml
Ms
First three describe the orbital, the location of the electron
Last one describes the particular electron (spin number)
34
First Quantum Number, the principal quantum number
it is n (energy level)
n=1,2,3… (natural numbers, starting at 1)
N is the main factor in determining energy, it is not the only factor
Average distance from the nucleus (electron can literally be anywhere in orbital but mostly, on average, it stays at the n distance)
Higher n is more energy
35
Second Quantum Number, the secondary quantum number
it is l (orbital shape)
L (lowercase, cursive)= 0,..., (n-1) (whole numbers)
The “l” value represents a specific orbital shape/sub shell/ sub level
L = 0 → s orbital → sphere
L = 1 → p orbital → figure 8 or dumbell
L = 2 → d orbital → clover shape
L = 3 → f orbital → look up picture…
Possible values of L depends on which energy level you are looking at
Secondary factor of energy
Shape of orbital
Higher l is more energy
36
Which sub-levels (orbital shape) exists for the principal quantum numbers, 1-4?
Other levels do exist but there are the only orbital shapes that exist in ground state
37
Third Quantum Number, the magnetic quantum number
it is ml (orbital orientation)
Ml (lowercase m with lowercase l as subscript)= -l,..., l (integers)
Related to the orientation of the orbital in space relative to the other orbitals in the atom
Number of orbitals for that sublevel= number of orientation
___Letter equals orbital but don’t forget to say + and -!
38
interesting orbital energy thing
4s is lower in energy than 3d
Periodic table is organized in pdf where s can have 2 electrons therefore 2 columns, f can have 14 electrons therefore, 14 columns which is based on the ml number and its possible electrons per orbital
39
Fourth Quantum Number, the spin quantum number
ms (spin)
In each sublevel (specific orbital), there are (up to) two electrons spinning in opposite directions
Whether electron is spinning clockwise or counterclockwise
ms= -½ or +½
*electron is a charged species which creates a magnetic field
40
pauli exclusion principle
no two electrons can have the same set of 4 quantum numbers therefore only 2 electrons/ orbital
41
aufbau principle
electrons will fill lowest energy orbitals first
* lowest energy = ground state
42
1s^1
```
1 = level
s= sublevel
^1= superscript of number of electrons in sublevel
```
43
electron configuration and box diagram
the gap between the boxes show the differences in energy.
44
hund's rule
- in a set of degenerate (same electron) orbitals, you will start by half filling the orbitals with electrons of the same spin
45
shorthand notation
includes previous noble and gas place it in sir brackets (shows that it includes all electrons of that gas)
- can be used for box diagrams as well
46
exceptions to octet rule
Be will be stable with 4 electrons since
| B will be stable with 6 electrons
47
when drawing a charged lewis structure
don't forget to include formal charge/ square brackets and right amount of electrons
48
expanded octets
remaining electrons are added to central atom. They are added to atoms past Phosphorus
49
things that can explain polarity
- shape does not determine polarity
| - non polar and polar molecules may contain polar bonds/ lone pairs
50
orientation of bonds and polarity
polar bonds on non polar molecule, they are on opposite sides of molecule
polar bonds in polar molecule, they are on same side of molecule
51
orientation of lone pairs and polarity
lone pairs on non polar molecule, they are on opposite sides of molecule
lone pairs on polar molecule, they are on same side of molecule
52
non polar molecules are
more symmetrical
53
dipole moment is measured in
debyes or D
54
bond angle for tetrahedral shape
109.5
55
effect of lone pairs on angles in VSEPR
more repulsion and take up more space so pushes bonds downwards and decreases them by about 1 to 2 degrees
56
5 VSEPR
trigonal bipyramidal
57
6 VSEPR
octahedral
58
5 with one lone pair VSEPR
see-saw
59
5 with 2 lone pair VSEPR
bent
60
5 with 3 lone pair
linear
61
6 with one pair
square pyramidal
62
6 with 2 lone pair
square planar
63
resonance structure
- when more than one possible lewis structure exists (by rotating pi bond around or lone pairs)
- actual molecule is a hybrid of the lewis structures
Best: formal charge is on most electronegative atom
64
benzene ring
draw a hexagon with a circle inside
65
hybridization
- occurs when the electron configuration does not match what the atoms oxidation state says
- used to make bonds of equal energy level (pi bonds are higher energy)
- when excited state, don't forget to put a star
66
hybridization and electron domain
```
2 --> sp
3 --> sp2
4-->sp3
5--> sp3d
6-->sp4d2
```
67
pi bonds and hybridization
need at least 1 separate orbital with an electron inside of higher energy
68
london dispersion
- temporary dipole caused by another molecule coming close
| - larger molecules are more stronger since there is more change of uneven distribution of charge at a single moment
69
dipole dipole
- permanent dipole
| - more polar the molecule, stronger the forces
70
hydrogen bonding
- H and FON,
due to small size and large electronegativity difference
- lone pair is attracted to the slightly positive end of hydrogen
- must have a lone pair on molecule
71
melting and boiling point hydrogen bonding
high boiling point means large IMF since the bonds must be broken with tons of energy to boil (see picture)
boiling point increases as size increases due to dispersion
72
viscosity hydrogen bonding
higher viscosity is high IMF since the molecules are more attracted to each other like honey with lots of glucose vs h2o
73
surface tension hydrogen bonding
high IMF means high surface tension
| tension at surface of a liquid
74
force of attraction in ionic compound is
Called a crystal lattice
| Attraction due to electrostatic forces of negative and positively charged ions
75
Composition in ionic compound is
Network of ions
| Formula is the ratio, not the number of atoms
76
Strength - Coulomb’s Law
Coulomb’s law calculates the force of attraction based on charges
+1 ion vs +2 ion, +2 ion will have greater attraction
77
Brittle (ionic compound)
Ions lined up perfect in an X and O pattern of anion to cation
When you take a hammer to an ionic compound, you shift the X and O pattern to be OXXO and XOOX (vs XOXO before) which creates large repulsion and it breaks
78
Electrical Conductivity (ionic compound)
Ions are not mobile so they do not conduct electricity (electrons hop on mobile ions and then get to destination in water)
In solution, anions have extra electrons and are mobile so they conduct electricity
79
What is the metallic bond?
When you take a metal like sodium (solid), you are looking at a bunch of sodium ions together
They create a sea of electrons where they each get rid of valence electrons and then the electrons just float around and go where needed (delocalized)
Matrix of positive nuclei in a sea of valence electrons
80
What affects the strength of the metallic bond?
The number of delocalized electrons creates more valence electrons which makes more binding force so higher melting point???
81
Electrical conductivity (properties of metals)
Since electrons are mobile and delocalized, can conduct electricity
82
Malleability(properties of metals)
When you take a hammer to it, the positive nuclei simply shift down and the sea of electrons just surrounds in again
There is no repulsion so it does not break
83
Low Volatility
| properties of metals
The delocalization is very strong ??
84
alloy
Mixture of metals (sometimes can be non-metal)
| 2 metals are melted, then mixed, then cooled
85
Alloys vs metals
Different sized radius
Does allow/it is harder for nuclei to slide past another
Results in greater strength
Different metals add different properties
86
covalent network
What is it?
A matrix where covalent bond holds atoms together
Covalent is very strong
Diamond (all sp3, tetrahedral)
also glass, SiO2
87
properties of covalent network
Very hard/rigid
Extremely high melting point
Often in group 14 (carbon and silicon)
88
Specific to Graphite structure
Structure: a network of c but hybridized sp2
Bunch of hexagons with 3 bonds per carbon
Flat (c is trigonal planar)
Multiple layers of carbon
No covalent bond between layers so attraction of layers is very weak (London dispersion)
89
Specific to Graphite properties
Softer (layers can slide)
| Conducts electrons through movement of electrons through resonance
90
vapor pressure hydrogen bonding
is the pressure of air combined with the few molecules that tend to evaporate and go back into liquid form over any liquid
- lower IMF means it is easier to get into gas phase which increases the vapour pressure
91
expanded octet
for third period atoms when there are too many electrons
92
deficient octet
boron and beryllium due to electronegativity different. It is high enough to participate in covalent bonding
93
formal charge
charge more than it should be on the atom
