unit 1 Flashcards
includes all unit 1 power points and study guide information. also all practice problems (not math)
1
Q
use substance and element in a sentence that describes relation
A
an element is a pure substance. the other pure substance is a compound.
2
Q
use substnace and mixture in a sentence that describes relation
A
a mixture is composed of 1 or more pure substances combined in variable proportions
3
Q
is it easier to prove an unknown substance an element or compound
A
compound
4
Q
mixtures are classified by their
A
properties
5
Q
classification of seawater
A
mixture
6
Q
4 properties of metals
A
malleable
ductile
lustrous
good conductors
7
Q
4 properties of nonmetals
A
poor conductors
brittle as solids
dull
gases at room temp
8
Q
1 chemical property of metals
A
oxides react with water to form hydroxides
9
Q
1 chemical property of nonmetals
A
oxides react with water to form acids
10
Q
is CO2 organic or inorganic
A
inorganic
11
Q
is H3PO4 molecular or ionic
A
molecular
12
Q
rf=
A
distance spot moved/distance solvent moved
13
Q
difference in ideas of atom development between democritus and dalton
A
democritus: no experimental evidence to support idea
dalton: had experimental evidence
14
Q
what evidence led thompson to determine that cathode ray was beam of negative particles
A
attracted to positive CRT
repelled by negative cathode
15
Q
why did rutherford conclude that nucleus has a positive charge
A
alpha particles have a positive charge. a few deflected off gold foil, Rutherford concluded deflected thing had a positive charge
16
Q
problems with rutherfordś model
A
e- should lose energy and be pulled into nucleus
atom should collapse
electrons are acellerating charges
17
Q
why was planckś theory not accepted at first
A
behavior of waves and particles are seen differently and supported by experiments and math
18
Q
E=
A
hv
h(c/lamda)
h=planck’s number
v=frequency
c=speed of light
lamda=wavelength
19
Q
h (planck’s number)
A
6.626 x 10^-34 J x s^-1
20
Q
6 types of radiation in order of lowest to highest eneergy
A
infared
red
blue-green
blue
violet
uv
21
Q
contrast Bohr and debroglie electrons
A
bohr: fixed energy levels, didnt know why they existed
debroglie: theoretical foundation for fixed energy levels that involved looking at wave properties of particles
22
Q
how did Heisenberg and Schrodinger see electron differently
A
Heisenberg: complex equations, particle w quantum behavior
Schrodinger: mathematics, wave phenomenon
23
Q
contrast borh’s energy level and quantum mechanical model
A
bohr: e- follow fixed paths around nucleus
qmeo: quantized energy, impossible to find electrons, e- found in atomic orbital, shows probability of finding an electron
24
Q
n shows
A
principal quantum number
size
energy levels
25
l shows
shape
number of different shapes (sublevels)
26
m
orientation in space
27
lobes of orbital disappear at nucleus. what does this mean?
0% chance of finding an e-
28
why does the 4s sublevel have lower energy than 3d
energies get closer together as n inclreases
repulsive forces means n has higher energy
3d has 5 orbitals while 4s has 1
29
allotropes
elements that exhibit different groupings of atoms
elements that border metalloids
ex. diamond, graphite, carbon, phosphorous
30
how do you know if a compound is organic?
contains carbon
almost all have a C-H bond
31
how do you know if a compound is ionic?
contains a metal and nonmetal
32
acids
contains hydrogen ion(s)
33
bases
contains/produces hydroxide ions
34
salts
ionic compounds without hydroxide
35
colloid
mixture that has very small particles distributed evenly in another substance
homogenous
36
homogeneous mixture properties
appears same throughout
particles smaller than 1 um
solution-substances do not aggregate to from particles (solute-solvent)
37
heterogeneous
doesn't appear same throughout
particles larger than 1 um
one or more compounds visible
38
suspension
dispersed phase and continuous medium
heterogenous
39
elements are classified based on their
properties
40
compounds are classified based on their
composition
41
potassium fluoride classification
compound
42
carbon classification
element
43
classification of something that contains more than one substance
mixture
44
density separation
used to separate solids with different densities
one liquid is denser than the other in the mixture
less dense liquids will float to top
must be insoluble in each other
45
filtrate
liquid that passes through the filter paper
46
mobile phase
liquid/gas carries mixture in chromatography
47
stationary phase
used to place the mixture on
mobile phase travels up this
components travel up this at different rates based on their attractions/affinities for this
substances with a strong attraction for this will not travel far (vice-versa)
48
paper chromatography
stationary phase: paper
mobile phase: solvent
49
filtration
mixture poured through filter to separate
50
decanting
gradually pour off a liquid to separate them
51
centrifugation
spinning solutions around an axis at high speed to separate them
52
column chromatography
stationary phase: tube packed with specially treated resin beads
mixture poured into one end of tube
mixture travels through beads
components separate
each component collected at the other end of the tube at different times
Rf= time component exited tube/time required solvent to exit
53
distillation
uses different vapor pressures/boiling points of mixture to separate
54
how to separate solution of aq copper sulfate, wood shavings, and iron filings?
use a magnet to attract iron filings
use filtration to separate wood and copper sulfate
distill copper sulfate from water
55
what's the difference between decomposing compounds and separating mixtures?
decomposing is chemical
separating mixtures is physical
56
explain chromatography
mobile phase carries each component of a mixture at a different rate through the stationary phase
each component travels at its own rate though the stationary phase based on its attraction to each phase
57
problems of distillation and how to fix it
liquids evaporate before they reach their boiling point
fix: distill the liquid several times to successfully remove the liquid with the higher boiling point
58
does chromatography or distillation require more energy?
distillation requires heat energy
chromatography relies on phase attractions
59
democritus
-proposed atoms
-no experimental evidence
60
law of conservation of mass
matter cant be created/destroyed
61
law of definite proportions
all samples in a given compound have the same proportions of their constituent elements
62
law of multiple proportions
Dalton reasoned that when 2 elements form 2 compounds, the products can be expressed as a ratio of whole numbers
63
atomic mass units are based on
the size of that atom compared to a hydrogen atom
64
Dalton's theory
all matter is composed of atoms which can't be created or destroyed
atoms of the same element are identical
atoms of different elements combine in simple, whole number ratios to form compounds
atoms are separated, rearranged, and recombined in chemical reactions to form new compounds
65
consistent with Dalton's theory?
sulfur and oxygen have the same mass
no
66
consistent with Dalton's theory?
all cobalt atoms are identical
yes
67
consistent with Dalton's theory?
K and Cl will combine in a 1:1 ratio to form KCl
yes
68
consistent with Dalton's theory?
Pb can be converted into Au
no
69
William Crookes discovery
began cathode-ray studies
70
JJ Thompson experiment
exposed cathode ray to electric field
-beam attracted to positive plate in electric field
repeated the experiment with different metals and gases, got the same results
71
JJ Thompson conclusions
beam composed of negatively charged (e-) particles
amount of deflection a charged particles experiences in an electric field depends on the ratio of the particle's charge to its mass
larger ratio = greater deflection
atoms are electrically neutral and contain positive and negatively-charged (e-) particles
72
charge-to-mass ratio for cathode rays
-1.76x10^8 C/g
73
millikan's oil drop experiment
gravity balance (droplets falling) and force of electrical field on negatively charged droplets
-repelled by negative plate
-attracted to positive plate
74
charge of an electron
-1.60 x 10^-19 C
75
thompson's plum pudding model
-used H2 in modified cathode ray tube
-proposed e- are evenly distributed through a mass of positive matter, like plums (e-) distributed through dough (+) of plum pudding
76
Rutherford's experiments findings
almost all mass in an atom is in the nucleus
cloud of e- makes up most of the volume in an atom, contributes little to its mass
e- move rapidly, held by the attraction of the nucleus
77
atomic number =
number of protons
78
mass number =
sum of protons + neutrons
79
nucleons
protons and neutrons
80
atomic mass unit
exactly 1/12 of the mass of a carbon-12 atom
81
mass spectrometry
measures masses of atoms and relative abundances
separates particles according to mass
82
how does a mass spectrometer work?
-atoms are injected and vaporized
-vaporized atoms are ionized by e- beam
-removes e-
-ions accelerated into magnetic field
-ions experience force that bends trajectory
-lighter ions are bent more easily
83
what part of dalton's theory is now incorrect?
-atoms aren't indivisible
-not all atoms of the same element are exactly the same, they can differ in the number of neutrons (isotopes)
84
Why did Rutherford believe that the nucleus was so small compared to the size of the atom itself?
-Most alpha particles pass through the gold foil showing that most of the volume of the atom is empty space
-Only a few particles bounced back from the tiny, dense nucleus
85
Rutherford describes the results of his experiment with this phrase: “It was almost as incredible as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you”. To what parts of the gold foil experiment was he referring to when he used this analogy?
shell = alpha particles
tissue paper = gold foil
86
electromagnetic spectrum
wavelengths
87
amplitude
height of wave
88
wavelength
distance between adjacent crests or troughs
89
frequency
number of cycles that pass a point per unit of time
90
wavelength and frequency proportion
inversely
91
(wavelength)(frequency) =
c
92
c
speed of light
93
speed of light =
3.0 x 10^8 m/s
94
planck's quantum theory
-electromagnetic waves' energy is proportional to frequency, inversely proportional to wavelength
-energy can behave like particles
-energy is be absorbed/emitted in quanta (whole numbers)
95
quanta
discrete packets of energy
96
waves
disturbances that move through space
can pass through and interfere with each other
can have any value within a range
97
E =
hv
or
h (c / wavelength)
98
particles
definite boundaries
bounce off each other when collided
can only exist in certain whole-number quantities
99
photoelectric effect
einsteins theory that used quantum mechanics to explain:
-metals emit electrons when a light shines on their surface
-light is composed of photons
100
photons
discrete packets of energy
101
Bohr discoveries
energy is quantized
hydrogen emitted four colored lines
different elements had their own patterns
H only has certain allowed energy levels, each corresponds to a circular orbit of a fixed size
larger orbits = greater energy
as long as an electron moves into an allowed state, the electron doesn't radiate or absorb energy
an electron can only move from one allowed orbit to another if it absorbs/emits an amount of energy exactly equal to the energy difference between the orbits
102
emission spectrum
spectrum released by each element
103
Bohrs postulates
-quantize energies for e- in H
-keep Rutherford's model from collapsing
-explain emission spectrum of H
104
quantum number
each allowed energy state is given an integer number "n" with allowed values from 1-infinity
105
ground state
lowest energy orbit (n=1)
106
excited states
larger orbits
107
issues with Bohr's model
didn't work with atoms of more than one e-
did not know why energy levels existed
108
DeBroglie's ideas
-matter is wave-like
-foundation for theories of e-'s fixed energy levels
-a whole-number multiple of wavelengths must fit into radius of an atom
109
DeBroglie's conclusion
the only allowed orbits for e- are those whose size (and energy) allows for an e- wave to be maintained
in the atomic world, we can't tell the difference between waves and particles
110
Heisenberg's uncertainty principle
impossible to locate an e- or where its going
-trying to measure this will cause the e- to change significantly
111
Schrodinger's Equation purpose
predicts the future behavior of a dynamic system
112
orbitals
3D regions where there is a high probability of finding an electron
113
What does it mean when we say that something is “quantized”?
It can only be absorbed or emitted in discrete “packets” or quantities
114
Why was Planck’s theory not accepted by most physicists at first?
The behavior of waves and particles was seen as different and this was supported by experimental evidence and the mathematics of the time
115
According to Planck, could an amount of energy equal to 2.5hv be emitted by an object?
No, energy can only be emitted in whole-number quantities
116
What finally led to Planck’s theory being accepted?
The discovery of the photoelectric effect. Einstein wrote a paper that used quantum theory to explain the photoelectric effect
117
Describe the appearance of hydrogen’s emission spectrum
There are four lines forming a noncontinuous spectrum. The lines are red (lowest energy), blue-green, blue and violet (highest energy)
118
Briefly describe how electrons generate each visible line in hydrogen’s emission spectrum
Each line represents an excited electron falling from a higher energy level to a lower energy level (n=2). The further the electron falls, the higher the energy of the light emitted.
Red: electrons fall from n=3 to n=2
Blue-green: electrons fall from n=4 to n=2
Blue: electrons fall from n=5 to n=2
Violet: electrons fall from n=6 to n=2
119
Why doesn’t hydrogen’s visible emission spectrum exhibit a continuous spectrum?
An electron can only emit light when it moves from one allowed orbit to another. It emits an amount of light equal to the energy difference between the two orbits
120
Which electron transitions in the emission spectrum, generate lines in the UV region of the electromagnetic spectrum
n = 6 to n = 1
121
List the six types of radiation in order from lowest energy to highest energy.
infrared, red, blue-green, blue, violet, ultraviolet
122
What important concept concerning atomic structure was Bohr’s theory responsible for?
Concept of fixed atomic energy levels
123
Why did Bohr’s model need to be replaced?
It didn’t work for atoms with more than one electron
124
How were the Bohr and De Broglie pictures of the electron different?
Bohr knew that energy levels existed but didn’t know why
De Broglie proposed a theoretical foundation for fixed energy levels that involved looking at the wave properties of particles
125
How did Heisenberg and Schrodinger see the electron differently?
Heisienberg’s theory treated the electron as a particle with quantum behavior
Schrodinger’s theory describe the wave properties of electron using mathematical equations
126
According to quantum mechanics, how can “throwing dice” apply to electron behavior?
Both use probability. Just as we can calculate the probability of throwing a 1,2 3,4,5, or 6 on a dice, we can also calculate the probability of finding an electron in a region of three-dimensional space.
127
Contrast Bohr’s electron orbit with the quantum mechanical electron orbital
Bohr: electrons follow fixed paths around the nucleus (planetary model)
QMT: probability volume shows where we are likely to find an electron around the nucelus
128
Schrodinger's Equation conclusion
the higher the energy, the greater the number and types of orbitals present
129
quantum numbers
specifies the characteristics of the orbitals and electrons
130
s orbital shape
spherical
131
p orbital shape
dumbell
132
principal quantum number (n)
Main energy level occupied by electron. Indicates
-relative size of orbital
-allowed energy states for the electron
Values are positive whole # integers (1,2,3,…)
As it increases, the electron:
-has more energy
-is farther from nucleus
Total number of orbitals in an energy level = n2
133
Hydrogen and n
ground state electron is in n=1
Spherical cloud with nucleus at center
Density of cloud is greatest near nucleus and decreases further from nucleus
1s orbital: spherical volume of space with a 90% chance of finding the hydrogen electron
134
to calculate number of orbitals:
n^2
135
sublevels determine
shape of the orbital
136
angular momentum quantum number (l)
ntegral values from 0 to (n-1) for each value of n
Ex: for n=1, only possible value is 0
Ex: for n=2, possible values of l are 0 and 1
Value of l is given by letters s, p, d , f
s=0
p=1
d=2
f=3
137
magnetic quantum number (ml)
orientation of the orbital in space around nucleus
Number of possible orientations =number of individual orbitals in the sublevel (s=1, p=3, d=5 etc.)
Values from +l to –l. Ex:
s sublevel: ml = 0 (1 orbital)
p sublevel: ml = -1, 0, +1 (3 orbitals)
d sublevel: ml = -2, -1, 0, +1, +2 (5 orbitals)
Each orbital holds maximum of 2 electrons
138
node
volume in space with 0 probability of finding an e-
139
Pauli Exclusion Principle
no two electrons in the same atom can have be described the same set of quantum numbers.
Two electrons in the same atomic orbital will have
the same first three quantum numbers.
opposite spins so their fourth quantum numbers are different.
The total number of orbitals existing in any energy level, n is equal to n2.
If two electrons can occupy each orbital, the maximum number of electrons that can exist in any energy level, n, is given by 2n2.
140
spin quantum number (Ms)
Spin state of the electron-magnetic field
only 2 possible directions (values): + ½ and - ½
paired electrons must have opposite spins
141
in hydrogen sublevels, all have
equal energy
142
in multi-electron atoms, evergies of sublevels are
different due to repulsive forces
143
energies get _ together as n inclreases
closer
144
repulsive forces cause some sublevels with smaller orbitals to have _ energies than larger orbitals
higher
145
State what each of the first three quantum numbers tell us about atomic orbitals
n: principal quantum number-relative size of the orbital and allowed energy states for the electron
l: number of different orbital shapes (sublevels) in an energy level
ml : orientation in space of an orbital
146
What is the difference between a 1s and a 2s orbital? What does that difference indicate about an electron possessing energy equal to n=2, compared to n=1?
2s orbital is larger and further away from the nucleus than the 1s orbital
An electron in n=2 has more energy than an electron in n=2
147
What are the two differences between a 2px and a 3py orbital?
The 3py orbital is located further from the nucleus and has a different orientation in space when compared with a 2px orbital
148
The lobes of a p-orbital disappear at the nucleus. What does this tell us about electrons in p-orbitals?
There is zero probability of finding an electron in the nucleus in p-orbitals
149
How many different orbitals are available to an excited hydrogen atom in the fourth energy level?
16
150
What does the fourth quantum number tell us about electrons?
the spin of the electron-up or down (+1/2 or -1/2)
151
For n=3, determine the number of allowed sublevels (different values of l) and give the designation of each
Value of l = 0 to n-1. For n=3, the value of l is 0-2.
There can be up to 3 sublevels in n=3 designated as: l=0 (s), l=1 (p), l= 2 (d)
152
Why can’t two electrons in the same orbital have the same four quantum numbers?
When two electrons occupy the same orbital, they must have opposite spins so ms for one electron is -1/2 and for the other electron is +1/2
153
What is the maximum number of electrons that can exist in the energy levels n=1 through n=4?
n=1, two electrons n=2, 8 electrons n=3, 18 electrons
n=4, 32 electrons
Total = 60 electrons
154
Explain why the 4s sublevel has a lower energy than the 3d sublevel
The 3d sublevel has 5 orbitals while the 4s only has one orbital.
Electron-electron repulsions in the five 3d orbitals give these orbitals a higher energy than the repulsions between the two electrons in the 4s orbital.
155
Aufbau principle
when filling orbitals, lowest energy orbitals are always filled first
156
electron configuration
format n (energy level) l (sublevel) #e- (number of electrons)
ex. H is 1s1
157
Hund's rule
when orbitals of equal energy are filled, most stable configuration has the maximum number of unpaired electrons with the same spin
158
diamagnetic
not magnetic, all electrons paired
159
paramagnetic
attracted to magnetic field. One or more unpaired electrons
160
how to draw orbital diagrams
Electrons –arrows point up or down to show opposite spins
Boxes or lines-orbitals within a sublevel
161
noble gas core condensed electron configuration
Symbol of nearest noble gas of lower atomic number
162
outer shell electrons condensed electron configuration
electron configuration for electrons after noble gas core
ex. Na is [Ne] 3s1
163
ion electron configuration
write normal configuration, then adjust by adding/removing electrons based on the ion
164
Silver has the condensed electron configuration [Kr] 5s1 4d10. This is not the configuration expected based on electron filling (using the diagonal rule). Write the expected electron configuration. Explain the anomalous behavior.
[Kr] 5s2 4d 9
A filled d-sublevel is lower energy (more stable) than a partially-filled d-sublevel
When Ag forms an ion it does so by losing a 5s electron rather than a 4d electron. This leaves it with a filled 4d sublevel
