Chem 1010 Final Flashcards
(135 cards)
1
Q
What is the limiting reagent?
A
The reactant that is completely consumed in a reaction, limiting the amount of product formed
2
Q
Theoretical yield vs actual yield
A
ty: amount of product that would be made in principle. based on the amount of limiting reactant
ay: how much product was actually procuded
3
Q
what is energy?
A
the capacity to do work, done when a force causes an object to move
4
Q
What are the three types of energy?
A
kinetic, potential, radiant
5
Q
potential energy
A
energy an object has by virtue of its position
6
Q
coulomb potential energy
A
potential energy that results from conservative coulomb forces and is associated with the configuration of a particular set of point charges within a defined system.
7
Q
electromagnetic energy
A
energy carried through space by the electromagnetic field
8
Q
how is potential energy affected when two opposite charges move closer to each other?
A
potential energy is negative and decreases, this is a more stable atom
9
Q
how is potential energy affected when two atoms of the same charge are brought close together?
A
the potential energy is positive and increases, this atom is less stable.
10
Q
how is potential energy affected by larger particles
A
repulsion (like charges) or attraction (opposite charges) increases.
11
Q
how is potential energy affected when particles are brought closer together?
A
attraction ( opposite charges) or repulsion (like charges) increases
12
Q
what is frequency
A
the number of crests/troughs that pass at a given point per unit time (Hz)
13
Q
what is the formula that relates the speed of light, frequency, and wavelength
A
c=v(lamda)
14
Q
constructive interference
A
waves collide in phase
15
Q
destructive interference
A
waves collide out of phase
16
Q
what happens when waves diffract through 2 slits
A
they interfere to produce a pattern of reinforcement/cancellation
17
Q
how does the angle of refraction change with wavelength
A
longer wavelengths refract more than shorter
18
Q
formula for potential energy
A
e=(1/2)mv^2
19
Q
ionization energy
A
the amount of energy required to remove the most loosely held electron in an atom
20
Q
what are photons
A
radiation that is absorbed or emitted in small packets
21
Q
what is the formula for the energy contained in one photon
A
E=hv=(hc)/v
22
Q
what is the photoelectric effect
A
light or photons with sufficient energy can dislodge an electron from a metal
23
Q
threshold fequency
A
the frequency of light which can remove an electron
24
Q
what must be absorbed for an electron to move from a lower to high energy level
A
a photon of energy (delta)E
25
formula for energy of the nth level
E=(-Rhc)/h^2
- <0 is an attractive force, lower energy, closer to nucleus
- n=1 is ground states
- n>1 is excited states
26
de brogle formula
(lamda)=h/mv
27
What are quantum numbers
a combination of 4 letters, unique to each electron, that describe the electrons location and properties
28
What is the Pauli Exclusion Principle?
states that no two electrons in an atom can have the same 4 quantum numbers
29
what are the 4 quantum numbers
```
n = principle quantum number
l = angular momentum quantum number
ml = magnetic quantum number
Ms = spin quantum number
```
30
what does the principle quantum number, n, represent
the size and energy of the orbital
n=1 is the smaller, e closest to the nucleus (lowest energy)
value of n denotes shell or level
31
what does the angular momentum quantum number, l, represent
```
the shape of the orbital
from 0 - n-1
0 - s orbital
1 - p orbital
2 - d orbital
3 - f orbital
```
32
what does the magnetic quantum number, ml, represent
specifies the orientation of the orbital
number of ml values for a given sub-shell is the number of possible orientations
-l,...,0,..., l
33
what does the spin quantum number, ms, represent
the spin of the electron. either +1/2 (up) or -1/2 (down)
34
probability density
probability per unit volume of finding an electron at a point in space surrounding a nucleus
35
nodes
places of very low electron probability
36
aufbau process
building up an element one electron at a time using these rules
1. e occupy orbitals in a way that minimizes energy of the atom
2. pauli exclusion principle
3. hunds rule
4. unpaired electrons have the same spin
37
hund's rule
when orbitals of identical energy are available electrons will occupy orbitals singly, to minimize repulsion
38
what is shielding
the blocking of valence shell electron attraction by the nucleus, due to the presence of inner-shell electrons. electrons in an s orbital can shield p electrons at the same energy level because of the spherical shape of the s orbital.
electron - electron repulsion "cancels" some of the attraction between nucleus and outer electrons.
39
formula for effective nuclear charge
Zeff= # of protons in nucleus - charge shielded by core electrons
40
what is penetration (in reference to shielding)
when outer electrons experience less of an attraction of the nucleus due to how well the outer electrons are shielded from the nucleus by the core electrons.
41
where are electrons removed from when cations are formed?
the highest n orbital
42
where do electrons go when anions are formed
the lowest available n orbital
43
what two things are affected by valence electrons
bonding and reactivity
44
what does inert means
unreactive
45
which columns on the periodic table are very reactive and why
```
column 1 (alkalis), column 7 (halogens)
- both are one electron away from a full valence shell
```
46
define a non-bonding atomic radius (van der walls atomic radius)
half the distance separating 2 nuclei during a collision in a gas phase, or between to adjacent nuclei in a solid
47
define bonding (covalent) atomic radius
half the distance between 2 nuclei of bonded atoms
48
what is the periodic trend for effective nuclear charge and atomic size?
they decrease across a row from left to right
| increase down a column (more electrons in a bigger shell)
49
how do transition metals vary from the trend for atomic size
they don't follow the trend for a size across a row. their size is constant across a row because the 3d orbitals are on average slightly closer to the nucleus than the 4s orbitals
50
are cations and anions smaller or larger than their parent atoms
cations are smaller, anions are larger
51
what is the size trend for an isoelectric series
ionic size decreases from left to right, because of differences in the effective nuclear charge
52
what is the trend for ionization energy
increases across a row (Zeff increases)
| increases up a column (atomic radii decreases, electrons more tightly held)
53
what is the exception to the ionization energy trend between the third and fourth column
a half filled p subshell is more stable than a p subshell with 4 electrons, so the ionization energy for those with half filled subshells (column 3) are higher.
54
what is electron affinity
the energy change when a gaseous atom gains an electron
55
what is the trend for electron affinity
increases up a group (added electrons are closer to nucleus, more stable)
increases across a row (higher Zeff)
56
which types of elements have no affinity
alkaline earth metals
noble gases
nitrogen atoms
57
what is magnetism
a force that can attract/repel objects that are made of magnetic material, behaviour in a magnetic field is explained by an atoms electron configuration
58
what creates a magnetic field
moving charges, so spinning electrons create a magnetic field, acting as a magnet
59
what is a paramagnetic atom
an atom with unpaired electrons
| they are strongly attracted to magnetic fields
60
what happens when all electrons are paired in an atom (in reference to magnetism) and what is an atom like this called?
individual magnetic effects cancel out. the two spins in each orbital must be opposite each other.
a diamagnetic atom
61
describe/explain metallic character
- metals have free electrons so they can conduct current
- the best indicator of metallic character is ionization energy
- metals have low ionization energies, so they form cations easily
62
what is the trend for metallic energy
increases down a group
decreases from left to right across a row
- hydrogen is non-metallic because it has a very large ionization energy
63
when do chemical bonds from
when two atoms have a lower potential energy together than they do separately.
64
octet rule
tendency for bonded atoms to transfer or share electrons in order to have an octet (noble gas configuration)
65
ionic bonds
transfer of electrons between two atoms
66
what holds together ionic bonds
large electrostatic attractions
67
what is lattice energy
energy released when ions are brought together to form a solid
68
what are some characteristics of ionic solids
- entire solid is ionic
- not held together by bonds between specific pairs
- all cations interact with all anions (attraction)
- all cation-cation, all anion-anion (repulsion)
69
outline the steps for drawing lewis structures for molecular compounds
1. write out correct skeletal structure
2. calculate total number of valence electrons
3. distribute electrons among atoms
4. form double or triple bonds if necessary
70
characteristics of double bonds
stronger than single bonds
shorter than single bonds
all electrons paired (even)
71
characteristics of triple bonds
strongest and shortest
tend to be un-reactive
all paired electrons
72
characteristics of pure covalent bonds
electrons shared equally between atoms
| only exits between identical atoms
73
characteristics of two non-identical atoms (referring to bond polarity)
electrons are shared unequally
| polar covalent bonds
74
electronegativity
ability of an atom to attract electrons to itself in a chemical bond
75
what are the electronegative values for non-polar covalent bonds
EN = 0-0.40
76
what are the electronegative values for polar covalent bonds
EN = 0.40-2.0
77
what are the electronegative values for ionic bonds
EN>2.0
78
Formal Charge
charge an atom would have if all bonding electrons are shared equally with others
formula = # of valence electrons - # of non-bonding electrons - (1/2) # of bonding electrons
79
where is the ideal place for any non-zero formal charges
on the most electronegative atom
80
what are free radicals
molecules with an odd amount of valence electrons
81
what atoms can form incomplete octets?
most often B or Be
82
what atoms can form expanded octets
S, P, I, atoms with d subshell, extra electrons can go there
83
VSEPR theory
valence-shell electron-pair repulsion
| predicts 3D shape and angles between nuclei
84
VSEPR rule (in reference to electron pair placement)
since electrons repel each other, electron pairs in the valence shell of the central atom seek to move as far apart as possible
85
what indicates the geometry of the molecule
the number of electron groups (lone pairs, single, double, triple bonds, lone electrons)
86
molecular geometry
arrangement of atoms around central atom
87
electron geometry
arrangement of electron groups around central atom
88
order of repulsion strength
lone pair - lone pair > lone pair - bond pair > bond pair - bond pair
89
where are lone pairs located for trigonal bypiramidal molecules
in equatorial positions
90
where are lone pairs located in octahedral molecules (when there are an even number of them)
opposite of each other
91
Valence bond theory
to say that an electron occupies an orbital is to describe a region of space in which electron density is the highest
covalent bond is the result of two unpaired electrons of opposite spin occupying an orbital formed by overlap of atomic orbitals.
92
sigma bond
a type of covalent bond that results from the head-to-head overlapping of atomic orbitals
93
pi bonds
covalent bonds that result from the lateral overlap of two atomic orbitals
94
steps for choosing hybridization scheme
1. use VSPER to predict electron geometry of atom
2. select hybrid orbitals that correspond to electron geometry
3. assume that terminal atoms are always un-hybridized
95
what kind of bonds are contained in a double bond
1 pi and 1 sigma bond
96
are sigma or pi bonds stronger
sigma bonds are stronger due to larger orbital overlap
97
what action do sigma bonds allow that pi bonds to not?
twisting around the bond
98
Molecular orbital theory
combination of the valence atomic orbitals of all of the atoms in the molecule to form a new orbital belonging to the whole molecule
99
LCAO
linear combination of atomic orbitals
100
formula for bond order
(1/2)(# of bonding electrons - # of anti-bonding electrons)
- keep in mind fractional bond orders are possible
101
what do bond orders 1, 2, 3 and those less than or equal to 0 indicate
1 - single bond
2 - double bond
3 - triple bond
less than or equal to 0 - bond doesn't form
102
are 2s or 2p molecular orbits lower in energy
2s
103
what does the state of a substance depend on
- thermal energy (total kinetic energy)
| - intermolecular forces
104
what are the 4 types of intermolecular forces
1. london dispersion forces (LDF)
2. dipole - dipole forces
3. hydrogen bonding
4. ion-dipole forces
105
what are LDFs
temporary attractive force that results from free floating electrons forming temporary dipoles
- present in all atoms and molecules
- increases with increasing molar mass
- globular molecules have lower LDFs than geometric molecules
106
what are dipole - dipole forces
when polar molecules align so that the positive end of one is near the negative end of another
- only present in polar molecules
107
what is hydrogen bonding
type of dipole - dipole attraction as a result of the attractive force between a hydrogen atom and a very electronegative atom such as N, O, F
- essentially an exceptionally strong dipole - dipole force
108
what are ion-dipole forces
exists between an ion and a polar molecule
| - dominate in aqueous solutions of ionic compounds
109
rank the intermolecular forces in order of decreasing strength
ionic, ion dipole > h-bonding > dipole-dipole > LDF (note - LDFs increasing with molar mass)
110
what are amorphous solids
solids with no defined faces, lacking in order
111
what are crystalline solids
atoms, ions, molecules are ordered in well-defined arrangements
112
```
state the
# of atoms per cell
# of nearest neighbours
formula for edge length
for a square cubic unit cell
```
1, 6, l=2r
113
```
state the
# of atoms per cell
# of nearest neighbours
formula for edge length
for a body centered cubic unit cell
```
2, 8, l=4r/(root3)=2.3r
114
```
state the
# of atoms per cell
# of nearest neighbours
formula for edge length
for a face centered cubic unit cell
```
4, 12, l=2(root2)r=2.8r
115
what are the three types of crystalline solids
molecular, ionic, atomic
116
what are molecular solids
individual units are molecules held together by intermolecular forces
Characteristics: low boiling points, held together by IMFs
117
describe ionic solids
individual units are ions
formed as a result of electrostatic attraction between cations and anions
characteristics: very strong, tend to have high melting points, brittle, don't conduct electricity, soluble in water
118
describe atomic solids
composed of individual atoms
119
what are the three types of atomic solids
nonbonding, metallic, network covalent
120
describe network covalent atomic solids
atoms that compose the solid held together by covalent bonds
- very high melting points because of strong covalent bonds
121
describe nonbonding atomic solids
form due to weak dispersion forces between atoms of noble elements
- low melting and boiling points
122
describe metallic atomic solids
aka metals
composed of metal cations in a sea of electrons
- electrons are free to move (delocalized) and attracted to cations (metallic bond)
- conduct electricity
- malleable because when stress is applied electron can move to prevent repulsion
- have variable melting points
- solid at room temp (except Hg)
123
what are three characteristics of ion arrangements
coordination number is maximized
charge neutrality is maintained (affects coordination number)
they are properly accommodated according to size
124
how can nuclear chemistry be used in medical diagnoses
radioactive tracers measure organ functions
125
what are the 5 types of radioactive decay
1. alpha decay
2. beta decay
3. gamma ray emisssion
4. positron emission
5. electron capture
126
what is alpha decay
radioactive decay in which a helium nucleus is emitted
- have a low penetrating power
- very large ionizing power
127
what is penetrating power
the ability to penetrating matter
128
what is ionizing power
the ability of radiation to ionize atoms and molecules
129
what is beta decay
occurs when an unstable nucleus emits a high speed electron
- low ionizing power due to smaller ion size
- higher penetrating power
130
what is gamma ray emission
high energy electromagnetic radiation that often occurs in conjunction with alpa and beta decay
- high penetrating power
- low ionizing power
131
what is positron emission
a proton is converted to a neutron to emit a positron
132
what is electron capture
the capture by the nucleus of an electron from the atom (its own)
133
what is a half life
the time it takes for half of a sample of radioactive atoms to decay
134
what is the rate equation
ln(Nt/No) = -kt
135
what is the equation for the half like of a substance
t(1/2)=0.693/k
