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final exam review Flashcards

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1
Q

what are sp hybrid orbitals

A

orbitals formed by mixing one s and one p orbital of a central atom

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2
Q

what is hybridization

A

a theoretical process of orbital mixing to form hybrid orbitals

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3
Q

what does the formation of bond orbitals mean

A

the wave functions of the electrons are in phase

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4
Q

what is the basic principle of valence bond theory

A

a covalent bond forms when orbitals of two atoms overlap and a pair of electrons occupy the overlap region

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5
Q

what does greater electronegativity difference result in

A

larger partial charges and higher partial ionic character

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6
Q

what is ON equal to

A

the number of valence electrons minus the number of shared electrons plus the number of unshared electrons

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7
Q

what is an oxidation number

A

the charge that an atom would have if all of the atoms attached to it were removed along with the electron pairs it is sharing with other atoms

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8
Q

electronegativity is ________ related to atomic size. why?

A

inversely related. the nucleus of a smaller atoms is closer to the shared pair than the nucleus of a larger atom

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9
Q

in substances what are most atoms joined by

A

polar covalent bonds, partly ionic and partly covalent

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10
Q

what makes a good fuel

A

covalent compounds with weak bonds and relatively less O

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11
Q

what is a fuel

A

a material that reacts with atmospheric oxygen to release energy

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12
Q

what is delta H r related to

A

sigma delta H reactant bonds broken + sigma delta H product bonds formed or sigma BE reactant bonds broken minus sigma BE product bonds formed *these are only for gas phase reactions

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13
Q

are covalent substances generally good conductors

A

no, electrons are localized and no ions are present

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14
Q

do molecular covalent substances have high melting and boiling points

A

no they have low mp and no because they have strong bonding forces between atoms within the molecule and weak intermolecular forces between separate molecules

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15
Q

do network covalent solids have a high melting and boiling points? why?

A

yes due to the large amount of energy required to rearrange covalent bonds

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16
Q

the trend in bond lengths is __________ to the trend in bond energy

A

opposite

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17
Q

as we go left across a group what does bond length do

A

increases

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18
Q

what does a higher bond order result in

A

a shorter bond length and higher bond energy

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19
Q

strong bonds are ______ in energy while weaker bonds are __________ in energy

A

lower, higher

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20
Q

why is bond energy always positive

A

bond breakage is an endothermic process

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21
Q

what is electron affinity

A

the energy change accompanying the addition of 1 mol of electrons to 1 mol of gaseous atoms of ions

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22
Q

what does the magnetic quantum number prescribe

A

the 3-d orientation of the orbital in the space around the nucleus

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23
Q

what are orbitals

A

mathematical solutions to the schrödinger wave equation

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24
Q

what are orbitals

A

mathematical solutions to the schrödinger wave equation

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25
Q

what does a higher nuclear charge do

A

it increases nucleus-electron attractions, lowering sub shell energy and stabilizing the atom

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26
Q

what is special about hydrogen and it’s energy state

A

it’s the only atom whose energy state depends completely on the principal quantum number

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27
Q

what does the principal quantum number specify

A

the energy level (shell) of the H atom, the higher the n value the higher the energy level

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28
Q

what is the principal quantum number

A

a positive integer that indicates the relative size of the orbital and therefore the relative distance from the nucleus of the peak in the radial distribution plot

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29
Q

what is psi (trident)

A

the wave function

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30
Q

what is the angular momentum quantum number related to (L)

A

the shape of the orbital

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31
Q

what is the major consequence of the exclusion principle

A

an atomic orbital can hold a maximum of two electrons and they must have opposing spins

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32
Q

what does penetration do

A

it increases the average effective nuclear charge for its electron and it decreases the shielding of a 2s electron by the 1s electrons

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33
Q

what does each solution to shrodingers equation give

A

an energy state associated with a given atomic orbital

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34
Q

:(

A

:(

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35
Q

what do transition metals lose first s or d electrons

A

s

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36
Q

what species are paramagnetic and what do they do

A

they are species with 1 or more unpaired electrons; they are attracted by an external magnetic field

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37
Q

elements in the same group have similar _______ and similar patterns of _______

A

valence electron configurations and reactivity

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38
Q

what are diamagnetic species and what do they do

A

they are species with all electrons paired, they are slightly repelled by an external magnetic field

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39
Q

are cations larger or smaller than parent atoms

A

smaller

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40
Q

are anions larger or smaller than parent atoms

A

larger

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41
Q

what do main group s block metals do

A

lose all electrons with the highest n balue

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42
Q

where do non metals gain electrons

A

in the p orbitals of the highest n value

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43
Q

what is penetration

A

the process by which an outer electron moves through the region occupied by the outer electrons to spend part of its time closer to the nucleus

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44
Q

what is the exclusion principle

A

no two electrons in the same atom can have the same four quantum numbers

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45
Q

which electrons are removed first from main group metals p or s

A

p

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46
Q

what does metallic behavior do across a period and down a group

A

in increases from right to left and down

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47
Q

what do elements that have low IEs and less negative EAs tend to do

A

lose electrons easily but attract them weakly

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48
Q

t/f elements that have high IEs and highly negative EAs tend to lose electrons with difficulty and attract them strongly

A

T

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49
Q

what does ionization energy do down a group? across a period?

A

decreases, increases

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50
Q

why does IE increase the more electrons you pull away

A

you are pulling electrons from a more and more positive ion

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51
Q

t/f ie2> ie1

A

t ALWAYS

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52
Q

what is ionization energy

A

the energy required to completely remove 1 mol of electrons from 1 mol of gaseous atoms or ions

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53
Q

as period number increases what happens to zeff

A

it rises significantly

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54
Q

which element is the biggest? the smallest?

A

Fr, F

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55
Q

what happens as zeff increases

A

outer electrons are pulled closer to the nucleus

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56
Q

what is bond energy

A

the energy needed to overcome the attraction between the nuclei and shared electrons, it is defined as the standard enthalpy change for breaking the bond in x mol of gaseous molecules

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57
Q

what is bond order

A

number of electron pairs being shared by a given pair of atoms

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58
Q

where is is the greater electron density in a covalent bond

A

between the nuclei

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59
Q

what causes a covalent bond

A

the balance between nuclei attracting the electrons and the electrons and nuclei repelling each other

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60
Q

what is the main way atoms interact

A

sharing electrons

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61
Q

why do ionic compounds have high mp and bp

A

large amounts of energy are required to free ions from their positions and separate them

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62
Q

when do ionic compounds conduct electricity

A

when melted or dissociated where ions can move and carry a current

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63
Q

what are typical physical behaviours of ionic compounds and why do they have these behaviours

A

they are hard (they do not dent), brittle (crack without deforming), and they are rigid (do not bend). there are strong attractive forces holding ions in specific positions

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64
Q

what does delta H lattice =

A

(cation charge * anion charge)/(cation radius + anion radius)

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65
Q

lattice energy is __________ to electrostatic energy

A

directly proportional

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66
Q

t/f there are periodic trends in lattice energy

A

t

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67
Q

what is lattice energy

A

the enthalpy change that accompanies the reverse of the equation of formation

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68
Q

is energy absorbed or released during electron transfer

A

absorbed

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69
Q

what form of matter are ionic compounds typically

A

solid

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70
Q

do 100% ionic or covalent bonds exist?

A

no, gradations of each type of bond occur

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71
Q
A
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72
Q

are electrons in metallic bonding localized or delocalized

A

deloacalized, they move freely throughout the piece of metal

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73
Q

in metallic bonding what is the metal iron core comprised of

A

nucleus plus inner electrons

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74
Q

what happens during metallic bonding

A

the enormous number of atoms in a sample of metal “pool” their valence electrons into a “sea” of electrons between and around each metal ion core

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75
Q

how does covalent bonding occur

A

the nucleus of each atom attracts the valence electrons of the other, drawing the atoms together

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76
Q

where does covalent bonding commonly occur

A

between non-metal atoms

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77
Q

what occurs during ionic bonding

A

electron transfer

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78
Q

where does ionic bonding occur

A

between atoms with large differences in electronegativity (usually between groups 1/2 and 17/the top of 16

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79
Q

what are resonance structures

A

two or more lewis structures for a molecule that cannot be adequately depicted by a single structure

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80
Q

what is electron pair deloaclization

A

the process by which electron density is spread over several atoms rather than remaining between the two atoms

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81
Q

what is a resonance hybrid

A

the weighted average of the resonance structures of a molecule

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82
Q

what is formal charge

A

the hypothetical charge on an atom in a molecule or ion, equal to the number of valence electrons minus the sum of all the unshared and half of the shared electrons

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83
Q

is formal charge the same as oxidation number

A

no

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84
Q

what are the electron deficient atoms

A

Be or B

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85
Q

what are free radicals

A

molecular or atomic species that contain one or more unpaired electrons which typically make the species very reactive

86
Q

what do most free radicals have

A

1) a central atom from an odd numbered group
2) a lone electron (this makes them paramagnetic)

87
Q

what are expanded valence shells

A

valence shells can accommodate more than 8 electrons by using available d orbitals

88
Q

when can expanded valence shells occur

A

only with central non-metal atoms from period 3 or higher

89
Q

when does the bond angle not equal the ideal angle

A

when there are double or triple bonds, when there are different elements, or when there is a lone pair

90
Q

what does AXmEn represent

A

A- central atom
X- surrounding atom
E- non-bonding valence electron groups

91
Q

what is the difference between electron group arrangement and molecular shape

A

-electron group arrangement is defined by bonding and nonbonding electron groups
-molecular shape is defined but the relative positions of the nuclei which are connected by bonding groups only

92
Q

what are metals good conductors of

A

electricity and thermal energy

93
Q

why do metals dent and bend

A

the metal cations are able to slide past each other in the “electron-sea” and end up in new positions

94
Q

what are alloys

A

mixtures that have metallic properties and consist of solid phases of two or more pure elements, a solid-solid solution or distinct intermediate phases

95
Q

which atoms will have more stable AOs

A

atoms with greater effective nuclear charge

96
Q

what is the effect of s and p orbital mixing

A

lower energy of the sigma 2s and anti sigma 2s

97
Q

in general the higher bond order the _______ the bond

A

stronger

98
Q

what is molecular orbital bond order

A

number of electrons in bonding MOs minus the number of antibonding MOs divided by 2

99
Q

what do electrons in an antibonding molecular orbital do

A

destabilize the molecule

100
Q

what is an antibonding molecular orbital

A

formed when wave functions are subtracted from each other decreasing the electron density between the nuclei and leaving a node

101
Q

what is MO theory

A

molecular orbital theory; describes a molecule as a collection of nuclei and electrons in which the electrons occupy orbitals extending over the entire molecule

102
Q

where do pi bonds occur

A

side to side overlap

103
Q

where do sigma bonds occur

A

end to end overlap

104
Q

when does hybridization not apply

A

to large non metal hydrides

105
Q

what shape does sp3d2 apply to

A

octahedral

106
Q

what shape does sp3d apply to

A

trigonal bipyramidal

107
Q

what shape does sp3 apply to

A

tetrahedral

108
Q

what shape does sp2 apply to

A

trigonal planar

109
Q

what shape does sp apply to

A

linear

110
Q

what are properties of gasses

A

they are highly compressible, thermally expandable, they have low viscosity and density, and they are infinitely miscible

111
Q

what are properties of an ideal gas

A

they have, negligible particle volume, many particles in random motion, particles colliding with each other and container walls, particles experience non interparticle forces, and the average kinetic energy of particles is directly proportional to the temperature of the gas

112
Q

what is urms

A

root mean squared speed equals the square root of u bar squared

113
Q

what does u bar squared equal

A

3RT/M

114
Q

do gasses really move at typical speed in bulk

A

no, individual particles are moving in all different directions, they collide with each other and change direction which slows them down

115
Q

what does temperature measure

A

the degree of random motion in a gas

116
Q

does temp have meaning for one particle

A

no it is a statistical measure and thus only meaningful for large numbers of particles

117
Q

what happens when we have 0 temperature

A

all translational motion ceases, however rotation and vibration can occur

118
Q

what conditions do we see ideal and non ideal behaviour of gases

A

ideal: low p high t
non-ideal: high p low t

119
Q

what causes non-ideal gases?

A

interparticle attraction

120
Q

what does P =

A

nRT/(V-nb) - a(n/v)^2

121
Q

in the van der waals equation what is a proportional to and b proportional to

A

a- interparticle forces
b- particle volume

122
Q

what do particles look like in a gas?

A

particles are far apart and moving randomly

123
Q

what do particles look like in a solid

A

they are close together and in an ordered arrangement

124
Q

what do particles look like in a liquid

A

they are close together but randomly oriwnted

125
Q

when we go from s-l-g energy is _____
what about g-l-s?

A

s-l-g: absorbed
g-l-s: released

126
Q

what is sublimation

A

solid to gas - delta H subl

127
Q

what is vaporization

A

liquid to gas - delta H vap

128
Q

what is fusion

A

melting, solid to liquid, delta Hfus

129
Q

what is condensation

A

gas to liquid- negative delta H vap

130
Q

what is freezing

A

liquid to solid, negative delta H fus

131
Q

what is deposition

A

gas to solid, negative delta H subl

132
Q

what does the magnitude of delta H reflect

A

the strength of interparticle forces

133
Q

why is delta H fus much less than delta H vap in general

A

much more energy needs to be absorbed to disrupt the IMFs and to separate particles far apart going from liquid to gas

134
Q

when are phase changes reversible

A

in closed systems where they are equilibrium processes

135
Q

when are phase changes not reversible

A

in open systems

136
Q

what is the vapour pressure of a liquid

A

the constant pressure at the dynamic equilibrium where the rate of vaporization matches the rate of condensation

137
Q

what does vapour pressure depend on

A

the type of liquid and the temperature

138
Q

why does higher temperature increase vapour pressure

A

more particles move faster exceeding the threshold speed to escape into gas phase

139
Q

what do substances with weaker intermolecular forces in the liquid have

A

higher vapour pressures and low boiling points

140
Q

when does boiling occur

A

when the vapour pressure equals the external pressure

141
Q

what is the normal boiling point

A

the temperature at which the vapour pressure is 1 atm

142
Q

why does pressure have little effect on melting points

A

solids and liquids are nearly incompressible

143
Q

what is happening at the melting point

A

a dynamic equilibrium between solid and liquid phases

144
Q

when will a solid sublime

A

when IMFs are too weak to keep particles together when they leave the solid

145
Q

what is present at the triple point

A

solid liquid and gas

146
Q

what is a supercritical fluid

A

occurs starting at the critical point, it is when the densities of liquid and gas become equal- we cannot distinguish between them

147
Q

is a supercritical fluid a mixture of solid and gas

A

NO

148
Q

what do intermolecular forces influence

A

the physical properties of a substance

149
Q

where are there no molecules

A

in simple ionic solids and extended network solids

150
Q

what are interparticle forces primarily based on

A

electrostatic attractions

151
Q

what is generally the strongest IMF? the weakest?

A

strongest- ion dipole
weakest- london dispersion (induced dipole dipole)

152
Q

why are ion dipoles strong

A

they have a full ion charge involved

153
Q

where do ion dipole forces arise

A

between ions and nearby polar molecules

154
Q

where do dipole dipole forces arise

A

between polar molecules, where dipoles are aligned to maximize attractions

155
Q

where do hydrogen bonding forces occur

A

between molecules that have H bonded to NOF

156
Q

why are hydrogen bonds strong

A

1) the difference in electronegativity is particularly large for N-H O-H and F-H
2) hydrogen is a small atom which allows the close approach of other dipoles

157
Q

what produces an induced dipole

A

shifts in electric field strength often produced by shifts in electron density

158
Q

when do induced dipole induced dipole forces arise

A

between non polar molecules or single atoms

159
Q

what is polarizability

A

the ease with which the electron cloud can be distorted to create a charge separation

160
Q

what increases polarizability

A

greater number of electrons and larger molecule size

161
Q

which is more polarizable cations or anions

A

anions because they have more electrons

162
Q

polarizability _______ down a group and _______ across a period

A

increases down a group and decreases across a period

163
Q

what does surface tension refer to

A

the resistance of a liquid to increase its surface area

164
Q

what does high surface tension imply

A

strong intermolecular forces

165
Q

what does capillarity refer to

A

the rising of liquid in a tube

166
Q

what is wetting

A

the spreading of a liquid across a surface

167
Q

what are cohesive forces

A

the forces within a liquid

168
Q

what are adhesive forces

A

the forces between the liquid and the tube or surface

169
Q

what happens to a liquid in a tube when cohesive and adhesive forces are similar in strength and character? different?

A

similar: the liquid is drawn up the tube
different: the liquid is not attracted to the tube

170
Q

what is viscosity

A

the resistance of a liquid to flow

171
Q

what does a high viscosity imply

A

high IMFs

172
Q

what happens to viscosity of a motor when temperature increases

A

the viscosity decreases as HC carbon chains are disentangled

173
Q

in XW30 what does a higher X value imply

A

more viscous oil

174
Q

what is the solute

A

the minor component

175
Q

what is the solvent

A

major component

176
Q

what is a good guideline for solubility

A

“like dissolves like”

177
Q

which electrons are involved in forming compounds

A

valence electrons

178
Q

what are outer electrons

A

those in the highest energy shell, they spend most of their time furthest from the nucleus

179
Q

what are core electrons

A

those that an atom has in common with the previous noble gas and any completed transition series, they fill all lower-energy shells of an atom

180
Q

what are half filled and filled sub shells

A

unexpectedly stable

181
Q

which subshell is more staple 2s or 2p

A

2s, it penetrates closer to the nucleus

182
Q

what does a lower nuclear charge do

A

makes the electron easier to remove

183
Q

what does shielding do

A

it reduces the full nuclear charge to an effective nuclear charge

184
Q

what are the two types of shielding

A

1) shielding by other electrons in a given shell
2) shielding by electrons in inner shells

185
Q

what are the two types of shielding

A

1) shielding by other electrons in a given shell
2) shielding by electrons in inner shells

186
Q

how many possible orientations does a d orbital have

A

5

187
Q

what is a 90% probability contour

A

the shape that the electron is within 90% of the time

188
Q

how well can an electrons position be known

A

with a certain probability

189
Q

as _______ increases, the probability that outer electrons spend most of their time farther from the nucleus increases

A

size

190
Q

as we move down a group what happens

A

n dominates and increases

191
Q

why does shielding occur

A

each electron feels both attraction to the nucleus and repulsion from other electrons, those repulsions counteract the nuclear attraction to a degree

192
Q

what are lobes

A

regions where it is highly probable to find an electron

193
Q

can we assign a definite size to an atom

A

no, the probability of finding an electron anywhere away from the nucleus is never 0

194
Q

what is the trident squared

A

probability density, a measure of the probability of finding the electron in some tiny volume of the atom

195
Q

what does removing an electron require

A

energy to overcome electrostatic attraction, IE is always positive

196
Q

what is electronegativity

A

the ability of an atom in a bond to attract electron density towards itself

197
Q

what is black body radiation

A

light given off by a hot idealized object that absorbed all of the radiation incident on it

198
Q

what is the lowest possible energy of a d subshell

A

3d

199
Q

the energy of an atom is quantized meaning?

A

it occurs in fixed quantities rather than being continuous

200
Q

what is an electron density diagram

A

an electron-cloud depicition

201
Q

what is the p orbital

A

the dumbbell or peanut shaped orbital

202
Q

what does the fourth quantum number describe (ms)

A

spin

203
Q

by convention what is the spin assigned to the first electron orbital

A

+1/2

204
Q

what is the s orbital

A

the spherical shaped orbital

205
Q

where does probability peak at (electron location)

A

near, but not at the nucleus

206
Q

what is a quantum

A

the definite amount of energy an atoms absorbs/emits when changing energy levels

207
Q

what is an orbital

A

a mathematical function that describes the electrons matter-wave (the three-dimensional space in which the highest probability exists of finding an electron)

208
Q

why do electrons have certain allowable energies and frequencies

A

because electrons have wave-length motion in orbitals of fixed radii

209
Q

what does the magnetic quantum number prescribe

A

the 3-D orientation of the orbital in the space around the nucleus

210
Q

what is electron affinity

A

the energy change accompanying the addition of 1 mol of electrons to 1 mol of gaseous atoms or ions

211
Q

what is urms =

A

sqrt(3RT/M)

chem 103 (3 decks)

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