General Inorganic Chemistry (Done) Flashcards
1
Q
Dalton’s Atomic theory states that? (5)
A
- Elements are composed of tiny indivisible particles called atoms.
- Atoms of a given element are identical in size, mass, and other properties. Atoms of different elements differ in size, mass, and other properties.
- Atoms cannot be subdivided, created, or destroyed.
- Atoms of different elements can combine in simple whole number ratios to form chemical compounds.
- In chemical reactions, atoms are combined, separated, or rearranged.
2
Q
This law states that if 2 elements combine to form one compound, the ratio of the mass of an element to a fixed mass of the other element is a whole number or a simple fraction.
A
Law of Multiple Proportions
3
Q
This law states that a chemical reaction only involves separation, combination or rearrangement of atoms.
A
Law of Conservation of Mass
4
Q
This person pioneered the atomic theory.
A
John Dalton
5
Q
Father of Chemistry; discovered the law of conservation of mass
A
Antoine Lavoisier
6
Q
This law states that a pure compound is made up of elements in the same proportion by mass, regardless of its quantity or source.
A
Law of definite proportions
7
Q
Discovered the law of definite proportions
A
Joseph Louis Proust
8
Q
Invented the cathode ray tube
A
Sir William Crookes
9
Q
The cathode ray tube was derived from _ by _.
A
Geissler Tube, Heinrich Geissler
10
Q
Discovered the electron through the cathode ray tube experiment
A
Joseph John (JJ) Thomson
11
Q
Ratio of electric charge to mass of electron
A
-1.76 x 10 ^ 8 C/g
12
Q
_ discovered the definite charge of electron (and proton) through the _ (experiment)
A
Robert Andrews Millikan, Oil Drop Experiment
13
Q
Definite charge of proton and electron
A
+- 1.6 x 10 ^ -19
14
Q
Discovered x-rays
A
Wilhelm Konrad Roentgen
15
Q
First to discovered evidence of radioactivity
A
Antoine Henri Becquerel
16
Q
Discovered radioactivity in Uranium and Polonium and coined the term radioactivity
A
Maria Skolowdowska (Marie) Curie
17
Q
_ discovered proton through _
A
Ernest Rutherford, gold foil experiment
18
Q
Discoveries of Ernest Rutherford in the Gold Foil Experiment (2)
A
- Most of the atom is empty space.
- Positive charge of an atom is concentrated at the nucleus.
19
Q
Mass of proton is _ and it is _ x mass of electron
A
1.6726 x 10 ^ 24 g
20
Q
Discovered neutron and its mass
A
James Chadwick
21
Q
Refers to the # of protons in the nucleus
A
Atomic number (Z)
22
Q
Atom is _ when # of protons = # of electrons
A
Electrically neutral
23
Q
This refers to the sum of the number of protons and neutrons
A
Mass number (A)
24
Q
Elements of the same atomic number but different mass numbers; same number of protons but different number of neutrons
A
Isotopes
25
Molecules of the same charge or same # of electrons
Isoelectronic
26
Elements of different atomic number but same mass number
Isobar
27
Elements of the same number of neutrons but different number of protons
Isotone
28
Weighted average of atomic masses of isotopes based on percent abundance
Average atomic mass
29
What are the four quantum numbers?
Principal, azimuthal/angular momentum, magnetic and spin quantum number
30
This quantum number refers to the average distance of electrons from nucleus in a particular orbital; determines the size of the orbital
Principal quantum number (n)
31
This quantum number refers to the shape of the orbital
Azimuthal/angular momentum quantum number (ℓ)
32
This quantum number refers to the orientation of orbitals in space
Magnetic quantum number (m_ℓ)
33
This quantum number refers to the spin of the electrons
Spin quantum number (m_s)
34
Principal quantum number (n) has values of?
Natural/counting numbers or 1, 2, 3, 4, ...
35
Azimuthal/angular momentum quantum number (ℓ) has values of
n - 1 or 0, 1, 2, 3, ...
36
Magnetic quantum number (m_ℓ) has values of
-ℓ to ℓ or ...,-3, 2, 1, 0, 1, 2, 3, ...
37
Spin quantum number (m_s) has values of
+1/2 (CW) or -1/2 (CCW)
38
This principle states that orbitals must be filled up in increasing energy levels
Aufbau Principle
39
Principle that states no two electrons can have the same set of quantum numbers and an atomic orbital must contain maximum of 2 electrons with opposite spins
Pauli Exclusion Principle by Wolfgang Pauli
40
Arranged the elements in triads
Dobereiner
41
Arranged elements in atomic mass and fund that every 8th element has similar properties
Newlands
42
He is the father of the modern periodic table and together with _, they arranged the elements according to their recurring periodic properties and increasing atomic mass
Dmitri Mendeleev, Lothar Meyer
43
Discovered the relationship between element's atomic number and frequency of x-rays generated when bombarded with high energy electrons
Moseley
44
Atomic property that refers to the distance between nucleus and valence electrons
Atomic size/radius
45
Electrons on the outermost shell
Valence electron
46
Electrons closer to the nucleus and have no participating in chemical bonding
Core electrons
47
Atomic property that refers to the energy required to remove an electron from a gaseous atom in its ground (lowest energy) state
Ionization energy
48
Atomic property that refers to the change in energy when an electron is accepted by a gaseous atom to form an anion
Electron affinity
49
Atomic property that refers to the measure of ability of atom to attract towards itself a bonding electron
Electronegativity
50
Atomic property that refers to the actual charge plus the repulsive effect due to shielding approximated by the difference between atomic number and the number of shielding/core electrons
Effective nuclear charge
51
True or False: Valence electrons shield core electrons better than core electrons shield each other.
False. Core electrons shield valence electrons better than valence electrons shield each other.
52
Atomic properties that decrease from left to right (or across the periodic table) and increase from top to bottom (or down the group)
Atomic size, metallic property, reactivity
53
Atomic properties that increase from left to right (or across the periodic table) and decrease from top to bottom (or down the group)
Ionization energy, electron affinity, electronegativity
54
Trend of effective nuclear charge
Increasing from left to right and top to bottom
55
Chemical bond that refers to the attraction of the nucleus and the electron of 2 or more atoms which usually occurs between 2 non-metals
Covalent bond
56
Chemical bond that refers to the attraction of 2 opposite charged particles (metal and non-metal)
Ionic bond
57
Property of chemical bond that refers to the energy released when a bond is formed
Bond energy
58
Property of chemical bond that refers to the distance between nuclei of the atoms forming the bond
Bond length
59
Property of chemical bond that refers to the number of bonding//shared pairs of electrons between two atoms
Bond order
60
Theory of bonding that assumes that electrons occupy atomic orbitals of the invidual atoms and states that bonds are formed due to overlap of 2 atomic orbitals.
Valence bond theory
61
Theory of bonding that states that electrons in the bonding molecular orbital are greater than the electrons in the non-bonding
Molecular orbital theory
62
This refers to chemical elements that is found naturally on Earth essentially as a single nuclide (which may, or may not, be a stable nuclide).
Monotopic/mononuclidic
63
How many elements are there that are monotopic?
21 elements (Be, F, Na, Al, P, Sc, Mn, Co, As, Y, Nb, Rh, I, Cs, Pr, Tb, Ho, Tm, Au, Bi, Pa)
64
Name this type of chemical reaction: A + B => AB
Combination/Synthesis
65
Name this type of chemical reaction: AB => A+B
Decomposition
66
Name this type of chemical reaction: AB + C => A + BC
Single displacement/replacement
67
Name this type of chemical reaction: AB + CD => AC + BD
Double displacement/decomposition
68
Combination of metal and oxygen forms
Basic oxide/anhydride
69
Combination of non-metal and oxygen forms
Acidic oxide/anhydride
70
Combination of non-metal oxide and water forms
Acid
71
Combination of metal oxide and water forms
Alkali/base
72
Decomposition of metal carbonate forms
Metal oxide and CO2
73
Decomposition of metal nitrates forms
Metal nitrites and O2
74
Decomposition of metal bicarbonate forms
Metal bicarbonate, water and CO2
75
Decomposition of oxyhalides form
Metal halides and oxygen
76
Activity series of halogens
F2 > Cl2 > Br2 > I2
77
Property that depends on the ratio of solute to solvent
Colligative properties
78
What are the four colligative properties?
Boiling point elevation
Freezing point depression
Vapor pressure lowering
Osmotic Pressure
79
Formula of boiling point elevation?
T_B,solution - T_B,solvent =K_b (m)
where T_B refers to boiling point temperature, K_b refers to the ebulliouscopic constant and m refers to the molality
80
Formula of freezing point depression
T_F,solution - T_F,solvent =K_F (m)
where T_F refers to freezing point temperature, K_F refers to the cryoscopic constant and m refers to the molality
81
Formula of vapor pressure lowering/Raoult's law
P_solvent,initial - P_solution = X_solute P_solvent,initial
where P refers to the pressure and X refers to the mole fraction
82
Formula of osmotic pressure (π)
π = MRT
where M is molarity, R is the ideal gas constant and T is the temperature in [K]
83
Molality is approximately the same as molarity IF _.
Solution is diluted.
84
Solutions of the same osmotic pressure (equal concentration) are considered to be _.
Isotonic
85
What are the ebullioscopic and cryoscopic constants of water?
Ebullioscopic - 0.52
Cryoscopic - 1.86
86
Relation between equilibrium constants Kp and Kc
Kp = Kc (RT) ^(Δn)
where n is the change in moles of gaseous products and reactants
87
This principle states that a system in equilibrium when subjected to a stress will act in such a way to relieve the stress (or shift back to equilibrium).
Le Chatelier's Principle
88
Increase in concentration of reactant results in
Forward reaction
89
Increase in concentration of product results in
Backward reaction
90
Increasing the pressure of the reaction results in
Shifting towards less gaseous molecules
91
Increasing the temperature when the reaction is endothermic results in
Forward reaction
92
Increasing the temperature when the reaction is exothermic results in
Backward reaction
93
Increasing the catalyst results in
No change in equilibrium
94
True or false: Catalyst affects the rate of reaction but not the equilibrium of the reaction.
True
95
SI unit of activity
Becquerel
96
The unit Curie refers to _.
Rate of disintegration of 1 g of Ra
97
SI unit of radiation absorbed dose
Gray
98
Effect of beta particle on (a) mass number (b) atomic number
(a) no effect
(b) decrease by 1
99
Effect of positron on (a) mass number (b) atomic number
(a) no effect
(b) increase by 1
100
Effect of proton or a H nucleus on (a) mass number (b) atomic number
(a) and (b) increase by 1
101
Effect of neutron on (a) mass number (b) atomic number
(a) increase by 1
(b) no effect
102
Effect of gamma ray on (a) mass number (b) atomic number
None
103
Effect of alpha particle or He nucleus on (a) mass number (b) atomic number
(a) increase by 4
(b) increase by 2
104
Nuclear reactions are balanced by:
1. Conservation of mass #
2. Conversation of atomic #
105
Alpha particle is essentially
He nucleus
106
Beta particle is essentially
Electron
107
Beta particle is emitted when _.
n/p is higher than the zone of stability
108
It is a by product of alpha particle decay and it is a high energy proton
Gamma ray
109
It is the antimatter of electron emitted when n/p is less than zone of stability
Positron
110
Inner orbital is captured by the nucleus to increase n/p
Electron capture
111
This is the force of attraction between nucleons over a distance (~ 1 x 10 ^ (-15) m) which overcomes the electromagnetic forces over short distances
Strong nuclear force
112
What does nucleon refer to?
Proton and neutrons
113
Value of n/p of stable nuclide
1
114
Value of n/p of unstable nuclides
> 1
115
Magic numbers of number of protons or neutrons
2, 8, 20, 50, 82, 126
116
This refers to the loss of mass that shows up as amount of energy released during nuclear transformation
Nuclear binding energy (change in E) = (change in mass) x (speed of light)^2
117
This refers to the change in mass during a nuclear transformation and it is computed through:
Mass defect (Δm) is computed by taking the difference between the mass of the products and the mass of the reactants.
118
True or false: a lower binding energy per nucleon implies more mass is converted to pure energy to bind the nucleons, hence indicating that the nuclei is more stable.
False. A /higher/ binding energy per nucleon implies more mass is converted to pure energy to bind the nucleons, hence indicating that the nuclei is more stable.
119
Coined "atomism"
Leucippus and Democritus
120
Coined atomos
Democritus
121
Proposed the law of multiple proportions
John Dalton
122
Region in space where an electron is most likely to be found is called
Orbital
123
Substances with unpaired electrons are called
Paramagnetic
124
Substances with paired electrons are called
Diamagnetic
125
Quantum Theory is proposed by _
Max Planck
126
Quantum Theory states that
"When solids are heated, they emit electromagnetic radiation over a wide range of wavelengths."
127
True or false: the amount of radiant energy emitted by an object at a certain temperature depends on its wavelength.
True
128
Max Planch assumed that atoms and molecules could emit or absorb energy in discrete quantities called _
quanta/quantum
129
Proposed the photoelectric effect
Albert Einstein
130
Albert Einstein proposed the _
Photoelectric effect
131
_ explains that electrons are ejected from the surface of certain metals exposed to a light of at least a certain minimum frequency called _.
Photoelectric effect, threshold frequency
132
Albert Einstein suggested that a beam of light is really a stream of particles called _
Photons
133
Dual nature of light states that
Light is both a particle and a wave
134
True or false: Photons cannot knock off electrons from surfaces of metals.
False. Photons can knock off electrons from surfaces of metals provided they have sufficient energy, i.e. energy is greater than the energy that binds the electrons to the metal (arising from attractive forces).
135
Force binding the electron to the metal is called _
Work function
136
Excess energy when electron is removed from the metal is converted into _ of the electrons that broke free.
Kinetic energy
137
Emission spectra of atoms is attributed to
Bohr
138
Bohr is attributed to
The emission spectra of atoms
139
Emission spectra are shown either as _ or _.
Continuous spectra or line spectra of radiation emitted by substances.
140
True or false: The emission spectra of atoms in the gas phase show a continuous spread of wavelength in the visible region.
False. It is not continuous.
141
Hydrogen spectral series that starts with the first orbital
Lyman: UV
142
Hydrogen spectral series that starts with the second orbital
Balmer: UV-Vis
143
Hydrogen spectral series that starts with the third orbital
Paschen: IR
144
Hydrogen spectral series that starts with the fourth orbital
Brackett
145
Describes the distribution of electrons in hydrogen and other atoms (address)
Quantum numbers
146
The Quantum numbers were derived from the mathematical solution of the _ for the hydrogen atom
Schrodinger equation
147
These are the wave functions of electrons in an atom; solutions to the Schrodinger equation
Atomic orbitals
148
This states that the square of the wave function determines the probability distribution of an electron in space.
Born interpretation
149
This principle states that the most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins
Hund's Rule of Maximum Multiplicity
150
How to compute for an estimate of the effective nuclear charge?
{# of protons} - {# of electrons inside}
151
_ devised a method for calculating relative electronegativities of most elements.
Linus Pauling
152
Representation of covalent bonding in which shared electron pairs are shown either as lines or pairs of dots between two atoms and lone pairs are shown as pairs of dots on individual atoms
Lewis Structure
153
True or false: Both core and valence electrons are shown in Lewis structures.
False. Only valence electrons are shown in Lewis structures.
154
An atom other than hydrogen tneds to form bonds until it is surrounded by _ valence electrons and this is described by the _.
eight, octet rule
155
Fajan's rule states that:
1. Small positive ion favors covalency.
2. Large negative ion favors covalency.
3. Large changers on either ion or on both ions favor covalency.
4. If the positive ion does not have a noble gas configuration, polarization/covalency is favored.
156
This has lower energy, which means greater stability than the atomic orbitals from which it was formed.
Bonding molecular orbital
157
This has higher energy, which means lower stability than the atomic orbitals from which it was formed.
Antibonding molecular orbital
158
This accounts for the geometric arrangements of electron pairs around a central atom in terms of the electrostatic repulsion between electron pair.
Valance-shell electron repulsion (VSEPR) model
159
Arrange in terms of strength: LP-BP repulsion, BP-BP repulsion, LP-LP repulsion (LP = lone pair, BP = bonding pair)
LP-LP > LP-BP > BP-BP
160
It is a quantitative measure of the polarity of a bond
Dipole moment
161
These are diatomic molecules containing atoms of different elements that have dipole moments
Polar molecules
162
These are diatomic molecules containing atoms of the same element that do not have dipole moments
Nonpolar molecules
163
Colligative properties talk about solutions that are _, with concentrations that are _.
(relatively) dilute, <0.2 M
164
Due to _, the vapor pressure of the solution of a nonvolatile solute is always _ than that of the pure solvent.
Vapor pressure lowering, less
165
Temperature at which the vapor pressure of a liquid is equal to the external atmospheric pressure.
Boiling point
166
Presence of a non-volatile solute _ the boiling point.
increases/elevates
167
A solution with nonvolatile solute has a _ boiling point than that of the pure solvent.
Higher
168
A solution with non-volatile solute has a _ freezing point that than of the pure solvent.
Lower
169
Selective passage of solvent molecules through a porous membrane from a dilute solution to a more concentrated on.
Osmosis
170
This is the pressure required to stop osmosis
Osmotic Pressure
171
If two solutions have unequal osmotic pressures, the more concentrated solution is called
Hypertonic
172
If two solutions have unequal osmotic pressures, the less concentrated solution is called
Hypotonic
173
True or false: The total number of solute particles determine the colligative properties of a solution
True
174
This refers to the ratio of actual number of particles in a solution after disassociation and the number of formula units initially dissolve in solution
Van't Hoff factor (i)
175
Van't Hoff factor has a value of _ for organics and non-electrolytic solutions
1
176
True or false: the colligative properties of electrolyte solutions are usually larger than anticipated.
False. Colligative properties are usually smaller than anticipated due to the formation of ion pairs at higher concentrations.
177
Formula for degree of disassociation (α)
α = (i-1)/(v-1)
where v is the maximum number of particles and i is the van't Hoff factor
178
Instead of the equilibrium constant, the _ is obtained when the reaction has not reached equilibrium.
Reaction quotient (Qc)
179
Direction of reaction when Qc < Kc
Forward
180
Direction of reaction when Qc > Kc
Backward
181
Direction of reaction when there is addition of inert gas at constant pressure
Towards the most mole of gas
182
Direction of reaction when there is addition of inert gas at constant volume
No change in direction
183
Equilibrium between the same substance of different phases
Physical equilibrium
184
Equilibrium between different substances of similar phases
Homogeneous equilibrium
185
State where rate of forward reaction = rate of backward reaction
Chemical equilibrium
186
Equilibrium between different substances of different phases
Heterogeneous equilibrium
187
This is the study of reactions involving changes in atomic nuclei
Nuclear chemistry
188
Spontaneous emission of particles and/or radition
Radioactivity
189
Any element that spontaneously emits radiation is said to be
Radioactive
190
Three types of rays produced by radioactive decay
Alpha, beta and gamma rays
191
True or false: Neutrons can be accelerated in a particle accelerator.
False
192
Densest element
Osmium
193
The nucleus disintegrates and emits particles and/or radition when _.
repulsion > attraction
194
The nucleus is stable when
Repulsion < attraction
195
This is the principal factor that determines the stability of nucleus.
Neutron to proton ratio (n/p)
196
n/p value becomes _ than 1 as atomic number increases
greater (since more neutrons are needed to counteract the stronger coulombic repulsion from protons)
197
All isotopes of the elements with atomic numbers greater than _, as well as the elements _ and _ are radioactive.
83, technetium, promethium
198
Stable nuclei are located within the _
Belt of stability
199
True or false: most radioactive nuclei liu outside the belt of stability
True
200
Binding energies per nucleon are greatest for group _ elements
8B (iron, cobalt and nickel region)
201
It is more meaningful to use the _ when comparing stability of any two nuclei
Nuclear binding energy per nucleon (due to the fact that they have different numbers of nucleons)
202
Formula for binding energy per nucleon
B = {[# of neutrons x mass of protons] + [# of neutrons x mass of neutrons] - [mass of atom]} x 931.49 MeV
203
All radioactive decay obey _ order of kinetics.
First
204
Formula for rate of decay at time t
t = λN
where λ is the rate constant and N is the number of radioactive nuclei present at time t
205
Formula for half-life of radioactive decay is
t_(1/2) = ln(2) / λ
206
It is the measure of the activity of the source of radioactivity
Disintegration rate
207
It is the measure of the energy deposited in a medium by ionizing radiation. It is used in the calculation of dose uptake in living tissue in both radiation protection and radiology.
Absorbed dose.
208
of proton, neutron and electron of Deuterium
1,1,1
209
Most common radioactive trace used in SPECT?
Tc-99
210
Most commonly used radioactive tracer in medicine?
Mo-99
211
Radioisotope used in liver
Fe-59
212
Radioisotope in carbon dating
C-14
