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unit 1 Flashcards

includes all unit 1 power points and study guide information. also all practice problems (not math)

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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.

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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

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3
Q

is it easier to prove an unknown substance an element or compound

A

compound

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4
Q

mixtures are classified by their

A

properties

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5
Q

classification of seawater

A

mixture

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6
Q

4 properties of metals

A

malleable
ductile
lustrous
good conductors

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7
Q

4 properties of nonmetals

A

poor conductors
brittle as solids
dull
gases at room temp

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8
Q

1 chemical property of metals

A

oxides react with water to form hydroxides

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9
Q

1 chemical property of nonmetals

A

oxides react with water to form acids

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10
Q

is CO2 organic or inorganic

A

inorganic

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11
Q

is H3PO4 molecular or ionic

A

molecular

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12
Q

rf=

A

distance spot moved/distance solvent moved

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13
Q

difference in ideas of atom development between democritus and dalton

A

democritus: no experimental evidence to support idea

dalton: had experimental evidence

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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

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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

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16
Q

problems with rutherfordś model

A

e- should lose energy and be pulled into nucleus

atom should collapse

electrons are acellerating charges

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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

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18
Q

E=

A

hv
h(c/lamda)

h=planck’s number
v=frequency
c=speed of light
lamda=wavelength

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19
Q

h (planck’s number)

A

6.626 x 10^-34 J x s^-1

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20
Q

6 types of radiation in order of lowest to highest eneergy

A

infared
red
blue-green
blue
violet
uv

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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

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22
Q

how did Heisenberg and Schrodinger see electron differently

A

Heisenberg: complex equations, particle w quantum behavior

Schrodinger: mathematics, wave phenomenon

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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

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24
Q

n shows

A

principal quantum number
size
energy levels

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25
Q

l shows

A

shape
number of different shapes (sublevels)

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26
Q

m

A

orientation in space

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27
Q

lobes of orbital disappear at nucleus. what does this mean?

A

0% chance of finding an e-

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28
Q

why does the 4s sublevel have lower energy than 3d

A

energies get closer together as n inclreases
repulsive forces means n has higher energy
3d has 5 orbitals while 4s has 1

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29
Q

allotropes

A

elements that exhibit different groupings of atoms

elements that border metalloids

ex. diamond, graphite, carbon, phosphorous

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30
Q

how do you know if a compound is organic?

A

contains carbon

almost all have a C-H bond

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31
Q

how do you know if a compound is ionic?

A

contains a metal and nonmetal

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32
Q

acids

A

contains hydrogen ion(s)

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33
Q

bases

A

contains/produces hydroxide ions

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34
Q

salts

A

ionic compounds without hydroxide

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35
Q

colloid

A

mixture that has very small particles distributed evenly in another substance

homogenous

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36
Q

homogeneous mixture properties

A

appears same throughout
particles smaller than 1 um
solution-substances do not aggregate to from particles (solute-solvent)

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37
Q

heterogeneous

A

doesn’t appear same throughout
particles larger than 1 um
one or more compounds visible

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38
Q

suspension

A

dispersed phase and continuous medium

heterogenous

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39
Q

elements are classified based on their

A

properties

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40
Q

compounds are classified based on their

A

composition

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41
Q

potassium fluoride classification

A

compound

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42
Q

carbon classification

A

element

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43
Q

classification of something that contains more than one substance

A

mixture

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44
Q

density separation

A

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

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45
Q

filtrate

A

liquid that passes through the filter paper

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46
Q

mobile phase

A

liquid/gas carries mixture in chromatography

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47
Q

stationary phase

A

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)

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48
Q

paper chromatography

A

stationary phase: paper
mobile phase: solvent

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49
Q

filtration

A

mixture poured through filter to separate

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50
Q

decanting

A

gradually pour off a liquid to separate them

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51
Q

centrifugation

A

spinning solutions around an axis at high speed to separate them

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52
Q

column chromatography

A

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

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53
Q

distillation

A

uses different vapor pressures/boiling points of mixture to separate

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54
Q

how to separate solution of aq copper sulfate, wood shavings, and iron filings?

A

use a magnet to attract iron filings
use filtration to separate wood and copper sulfate
distill copper sulfate from water

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55
Q

what’s the difference between decomposing compounds and separating mixtures?

A

decomposing is chemical
separating mixtures is physical

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56
Q

explain chromatography

A

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

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57
Q

problems of distillation and how to fix it

A

liquids evaporate before they reach their boiling point

fix: distill the liquid several times to successfully remove the liquid with the higher boiling point

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58
Q

does chromatography or distillation require more energy?

A

distillation requires heat energy

chromatography relies on phase attractions

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59
Q

democritus

A

-proposed atoms
-no experimental evidence

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60
Q

law of conservation of mass

A

matter cant be created/destroyed

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61
Q

law of definite proportions

A

all samples in a given compound have the same proportions of their constituent elements

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62
Q

law of multiple proportions

A

Dalton reasoned that when 2 elements form 2 compounds, the products can be expressed as a ratio of whole numbers

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63
Q

atomic mass units are based on

A

the size of that atom compared to a hydrogen atom

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64
Q

Dalton’s theory

A

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

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65
Q

consistent with Dalton’s theory?
sulfur and oxygen have the same mass

A

no

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66
Q

consistent with Dalton’s theory?
all cobalt atoms are identical

A

yes

67
Q

consistent with Dalton’s theory?
K and Cl will combine in a 1:1 ratio to form KCl

A

yes

68
Q

consistent with Dalton’s theory?
Pb can be converted into Au

A

no

69
Q

William Crookes discovery

A

began cathode-ray studies

70
Q

JJ Thompson experiment

A

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
Q

JJ Thompson conclusions

A

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
Q

charge-to-mass ratio for cathode rays

A

-1.76x10^8 C/g

73
Q

millikan’s oil drop experiment

A

gravity balance (droplets falling) and force of electrical field on negatively charged droplets
-repelled by negative plate
-attracted to positive plate

74
Q

charge of an electron

A

-1.60 x 10^-19 C

75
Q

thompson’s plum pudding model

A

-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
Q

Rutherford’s experiments findings

A

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
Q

atomic number =

A

number of protons

78
Q

mass number =

A

sum of protons + neutrons

79
Q

nucleons

A

protons and neutrons

80
Q

atomic mass unit

A

exactly 1/12 of the mass of a carbon-12 atom

81
Q

mass spectrometry

A

measures masses of atoms and relative abundances

separates particles according to mass

82
Q

how does a mass spectrometer work?

A

-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
Q

what part of dalton’s theory is now incorrect?

A

-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
Q

Why did Rutherford believe that the nucleus was so small compared to the size of the atom itself?

A

-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
Q

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?

A

shell = alpha particles
tissue paper = gold foil

86
Q

electromagnetic spectrum

A

wavelengths

87
Q

amplitude

A

height of wave

88
Q

wavelength

A

distance between adjacent crests or troughs

89
Q

frequency

A

number of cycles that pass a point per unit of time

90
Q

wavelength and frequency proportion

A

inversely

91
Q

(wavelength)(frequency) =

A

c

92
Q

c

A

speed of light

93
Q

speed of light =

A

3.0 x 10^8 m/s

94
Q

planck’s quantum theory

A

-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
Q

quanta

A

discrete packets of energy

96
Q

waves

A

disturbances that move through space
can pass through and interfere with each other
can have any value within a range

97
Q

E =

A

hv
or
h (c / wavelength)

98
Q

particles

A

definite boundaries
bounce off each other when collided
can only exist in certain whole-number quantities

99
Q

photoelectric effect

A

einsteins theory that used quantum mechanics to explain:
-metals emit electrons when a light shines on their surface
-light is composed of photons

100
Q

photons

A

discrete packets of energy

101
Q

Bohr discoveries

A

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
Q

emission spectrum

A

spectrum released by each element

103
Q

Bohrs postulates

A

-quantize energies for e- in H
-keep Rutherford’s model from collapsing
-explain emission spectrum of H

104
Q

quantum number

A

each allowed energy state is given an integer number “n” with allowed values from 1-infinity

105
Q

ground state

A

lowest energy orbit (n=1)

106
Q

excited states

A

larger orbits

107
Q

issues with Bohr’s model

A

didn’t work with atoms of more than one e-

did not know why energy levels existed

108
Q

DeBroglie’s ideas

A

-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
Q

DeBroglie’s conclusion

A

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
Q

Heisenberg’s uncertainty principle

A

impossible to locate an e- or where its going

-trying to measure this will cause the e- to change significantly

111
Q

Schrodinger’s Equation purpose

A

predicts the future behavior of a dynamic system

112
Q

orbitals

A

3D regions where there is a high probability of finding an electron

113
Q

What does it mean when we say that something is “quantized”?

A

It can only be absorbed or emitted in discrete “packets” or quantities

114
Q

Why was Planck’s theory not accepted by most physicists at first?

A

The behavior of waves and particles was seen as different and this was supported by experimental evidence and the mathematics of the time

115
Q

According to Planck, could an amount of energy equal to 2.5hv be emitted by an object?

A

No, energy can only be emitted in whole-number quantities

116
Q

What finally led to Planck’s theory being accepted?

A

The discovery of the photoelectric effect. Einstein wrote a paper that used quantum theory to explain the photoelectric effect

117
Q

Describe the appearance of hydrogen’s emission spectrum

A

There are four lines forming a noncontinuous spectrum. The lines are red (lowest energy), blue-green, blue and violet (highest energy)

118
Q

Briefly describe how electrons generate each visible line in hydrogen’s emission spectrum

A

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
Q

Why doesn’t hydrogen’s visible emission spectrum exhibit a continuous spectrum?

A

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
Q

Which electron transitions in the emission spectrum, generate lines in the UV region of the electromagnetic spectrum

A

n = 6 to n = 1

121
Q

List the six types of radiation in order from lowest energy to highest energy.

A

infrared, red, blue-green, blue, violet, ultraviolet

122
Q

What important concept concerning atomic structure was Bohr’s theory responsible for?

A

Concept of fixed atomic energy levels

123
Q

Why did Bohr’s model need to be replaced?

A

It didn’t work for atoms with more than one electron

124
Q

How were the Bohr and De Broglie pictures of the electron different?

A

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
Q

How did Heisenberg and Schrodinger see the electron differently?

A

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
Q

According to quantum mechanics, how can “throwing dice” apply to electron behavior?

A

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
Q

Contrast Bohr’s electron orbit with the quantum mechanical electron orbital

A

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
Q

Schrodinger’s Equation conclusion

A

the higher the energy, the greater the number and types of orbitals present

129
Q

quantum numbers

A

specifies the characteristics of the orbitals and electrons

130
Q

s orbital shape

A

spherical

131
Q

p orbital shape

A

dumbell

132
Q

principal quantum number (n)

A

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
Q

Hydrogen and n

A

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
Q

to calculate number of orbitals:

A

n^2

135
Q

sublevels determine

A

shape of the orbital

136
Q

angular momentum quantum number (l)

A

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
Q

magnetic quantum number (ml)

A

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
Q

node

A

volume in space with 0 probability of finding an e-

139
Q

Pauli Exclusion Principle

A

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
Q

spin quantum number (Ms)

A

Spin state of the electron-magnetic field

only 2 possible directions (values): + ½ and - ½

paired electrons must have opposite spins

141
Q

in hydrogen sublevels, all have

A

equal energy

142
Q

in multi-electron atoms, evergies of sublevels are

A

different due to repulsive forces

143
Q

energies get _ together as n inclreases

A

closer

144
Q

repulsive forces cause some sublevels with smaller orbitals to have _ energies than larger orbitals

A

higher

145
Q

State what each of the first three quantum numbers tell us about atomic orbitals

A

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
Q

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?

A

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
Q

What are the two differences between a 2px and a 3py orbital?

A

The 3py orbital is located further from the nucleus and has a different orientation in space when compared with a 2px orbital

148
Q

The lobes of a p-orbital disappear at the nucleus. What does this tell us about electrons in p-orbitals?

A

There is zero probability of finding an electron in the nucleus in p-orbitals

149
Q

How many different orbitals are available to an excited hydrogen atom in the fourth energy level?

A

16

150
Q

What does the fourth quantum number tell us about electrons?

A

the spin of the electron-up or down (+1/2 or -1/2)

151
Q

For n=3, determine the number of allowed sublevels (different values of l) and give the designation of each

A

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
Q

Why can’t two electrons in the same orbital have the same four quantum numbers?

A

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
Q

What is the maximum number of electrons that can exist in the energy levels n=1 through n=4?

A

n=1, two electrons n=2, 8 electrons n=3, 18 electrons
n=4, 32 electrons
Total = 60 electrons

154
Q

Explain why the 4s sublevel has a lower energy than the 3d sublevel

A

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
Q

Aufbau principle

A

when filling orbitals, lowest energy orbitals are always filled first

156
Q

electron configuration

A

format n (energy level) l (sublevel) #e- (number of electrons)

ex. H is 1s1

157
Q

Hund’s rule

A

when orbitals of equal energy are filled, most stable configuration has the maximum number of unpaired electrons with the same spin

158
Q

diamagnetic

A

not magnetic, all electrons paired

159
Q

paramagnetic

A

attracted to magnetic field. One or more unpaired electrons

160
Q

how to draw orbital diagrams

A

Electrons –arrows point up or down to show opposite spins

Boxes or lines-orbitals within a sublevel

161
Q

noble gas core condensed electron configuration

A

Symbol of nearest noble gas of lower atomic number

162
Q

outer shell electrons condensed electron configuration

A

electron configuration for electrons after noble gas core

ex. Na is [Ne] 3s1

163
Q

ion electron configuration

A

write normal configuration, then adjust by adding/removing electrons based on the ion

164
Q

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.

A

[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

ap chem (11 decks)

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