Structure Of Matter Flashcards
1
Q
Antoine Lavoisier
A
Credited with the formulation of the law of conservation of matter in 1789
2
Q
Law of conservation of mass
A
In a chemical reaction, matter is niether created nor destroyed
Key concept to the formulation of modern atomic theory
3
Q
Joseph Proust
A
Credited with the formulation of the law of definite proportions in 1797
4
Q
Law of definite proportions
A
All samples of a given compound have the same proportions of their constituent elements
5
Q
Mass ratio
A
The number of each elements in the compound if the fewest elements number just one
Eg: nitrogen to carbon ratio of 4.7:1
6
Q
John Dalton
A
Credited with the formulation of the law of multiple proportions in 1804, and then the atomic theory in 1808
(Many other works too, should probobly check wikipedia for more)
7
Q
Law of multiple proportions
A
When two elements form two different compounds the masses of the elements can be described in a mass ratio
8
Q
Atomic theory
A
1) each element is composed of tiny, indestructable particles called atoms
2) all atoms of a given element have the same mass and other properties that make them unique from any other element
3) atoms combine in simple, whole number ratios to form compounds
4) atoms od one element cannot change into atoms of another element, but can change the way in which they are bonded to other atoms
9
Q
Atomic mass
A
The average mass of one mole (6.022x10^23) of each element
Given in amu (equals g/mol)
Not constant from one sample to another
Sometimes given as a range in which that mass could fall
10
Q
Physical determination of atomic mass
A
Sum of the products of each isotope’s decimal abundance, mass of each isotope
11
Q
Mass spectrometery
A
Method of separating atoms by their mass
Sample injected into cylinder, vaporized, ionized by electrons, accelerated through a magnetic field.
The change in their trajectory due to this magnetic force tells us their individual masses, lighter ions experience greater change
12
Q
Atomic number (z)
A
Tells 1) number of protons 2) number of electrons (when nuetral) 3) positive charge of the nucleus Calculated as z=A-N
13
Q
Neutron Count (N)
A
Number of neutrons in a given atom
Calculated as N=A-z
14
Q
Mass number (A)
A
Total number of nucleons (protons and neutrons) in an atom A=z+N Often written as [Element name]-A (eg: carbon-12)
15
Q
Isotopes
A
Atoms with the same number of protons but different neutrons
16
Q
Isotope notation
A
A over z and then the element name
Mass number is always written over the number of protons
17
Q
Natural abundance
A
Percentage breakdown of isotopes present in any particular sample of an element, varies by isotope
18
Q
Electron energy levels
A
Each row of the periodic table is a new energy level for electrons
The ‘unfilled level’ being the level containing valance electrons
19
Q
Electron configuration
A
Way of showing which electrons occupy which orbitals
Ground state= H 1s^1
20
Q
Pauli exclusion principal
A
No two electrons can ever have the same four quantum numbers
21
Q
Orbital
A
Any one of four areas that describe the entire distribution of where that electron might be located.
22
Q
P orbital
A
Electrons here occupy an any space within a peanut shape arround the nucleus
23
Q
D orbital
A
Electrons here occupy an any space within a flower shape arround the nucleus
24
Q
F orbital
A
Electrons here occupy an any space within a 3D-flower shape arround the nucleus
25
Quantum numbers
Values that describe the behavior of an individual electron within an atom
26
S orbital
Lowest energy
| Electrons here may occupy a sphere-like area around the nucleus
27
Principal quantum number (n)
Positive integer that represents electron energy level
28
Angular momentum quantum number(l)
Integer that represents the shape of the orbital
S=0
P=1
D=2
F=3
29
Magnetic quantum number (m sub-l)
Integer that represents orbital orientation
Exist on the interval (-l to +l)
l is the angular momentum quantum number
30
Spin quantum number (m sub-s)
Refers to the direction in which the electron spins
Either -1/2 or +1/2
31
Periodic property
A property of an atom predictable by its position on the periodic table
32
Electron configuration
A way of writing which electrons are in which orbital
'Ground State'- 1s^1
33
Dmitry Mendeleev
Russian chemist Credited with the arrangement of the periodic table based on the work of German chemist Julius Lothar Meyer
34
Orbital Diagram
Way of drawing the 'direction' of electron 'spin' within electron configuration
Little arrows going up and down
35
Pauli Exclusion Principal
No two electrons can have the same four quantum numbers within the same atom
36
Coulombs law
E=q1q2/r*1/(4(pi)*(8.85x10^-12))
The energy between two charged particles is equal to the product of their charges divided by the distance between them over 4pi*energy constant
1) energy decreases with distance
2) opposite charges attract one another, while particles of opposite charges repel
3) particles are attracted to one another based on the magnetude of the opposing charge
37
Shielding
The property of an electron to repel another electron
38
Effective Nuclear charge
Refers to the charge acting on a particular electron
The ionic charge of that atom if the electron did not exist
39
Penetration
The action of an electron changing energy levels due to an attraction to the Nucleus
40
Aufbau Principal
Only two electrons of opposite spins are allowed in each orbital
41
Hunds Rule
Electrons must fill these orbitals as single electrons before they can double up
42
Valance Electrons
Electrons that are exchanged during chemical bonding
43
Core electrons
Electrons that would compose an atom of the noble nearest two but less than the mass of the atom in question
44
Orbital Blocks of the periodic table
A way of demonstrating on the periodic table which orbitals are used in electron configuration
G1and2 = s
Transitions = d
Metaloids/nonmetals = p
Weird rows at the bottom = f
45
Noble gasses
Elements of the row furthest to the right of the periodic table
Rarely bond with any other element
Used in experiments in which a reaction is undesireable
46
Alkali Metals
Group 1 elements, highly reactive
One valence electron
Excelent reducing agents (decrease charge of a material by contributing an electron)
Result in violent reactions
47
Alkali Earth Metals
Group 2 elements
| Two valence electrons
48
Halogens
Group 7
7 valence electrons
Excelent oxideizing agents (take electrons from a material thereby increasing their charge)
49
Van der wals radius
The radius of an atom when it is not bonded to another atom
50
Covalent radius
Radius of an atom when it is bonded to another atom
51
Atomic radii
Lower left of the periodic table
52
Electronegativity
Ability for an atom to attract electrons
Opposite of metallic Character
Increases to the upper right of an atom
53
Metallic Character
Ability for an atom to loose an electron
Opposite of electronegativity
Increases to the lower left of the periodic table
54
Paramagnetism
Property of atoms to be attracted to an external magnetic field
An effect of an empty space in a low energy orbital
Usually attributed to silver
55
Diamagnetic
Property of atoms to be repelled by an external magnetic field
An effect of full energy orbitals
Usually attributed to zinc
56
Cation radius rule
Cations are nearly half the size of their nuetral atoms
57
Anion radius rule
Anions are much larger than their neutral atoms
58
Electron affinity
The energy change associated with gaining an electron in a gaseous state
Almost always negative
Increases (from negative... Approaches zero) to the upper left of the periodic table
59
Ionic bonds
Oppositely charges ions attract one another for an exceptionally low net charge
Electrons are tranfered in this bond
Electronegativity <.5
60
Covalent bonds
Two nonmetals bonding together
Form molecules
Atoms share their electrons
Electronegativity .5
61
Metallic bonds
Metals bond to an atom of the same element by pooling their vallence electrons (electron sea model)
No positive metal ions are attracted to the 'sea' of electrons
Get clarification on this
62
Macromolecular bonding
Smaller molecules bonded together in a polymer chain
63
Hydrogen bonding
Strong dipole-dipole bond between Hydrogen and either O, N, or F
64
Polar covalent bonds
High electronegativity difference between two atoms bonded together
Electronegativity > 1.7
65
Pi bond
Bond that forms betwwen two overlapping P orbitals
66
Peptide bond
Bond that forms between the amine end of one amino acid and the carboxylic end of another
67
Sigma bond
Bond resulting from an overlap of a p orbits and an s orbital
Commonly sp^2 hybridization
68
Valence bond theory
Advance bonding model in which electrons reside in quantum mechanicalorbits localized in individual atoms, a hybridized blend of standard atomic orbitals
A bond occurs when these orbitals overlap
69
Lewis theory
Simple model of of chemical bonding in which atoms are arranged to form 'octets' for each atom
70
Lewis structures
Drawings of atoms in which bonds are represented as dots to form octets
71
Octet rule
When bonded together, electrons must share as many electrons as it takes so that they each have ready access to 8 electrons
72
Lewis base
An atom, ion, or molecule that donates an electron pair
73
Lewis acid
An atom, ion, or molecule that accepts an electron pair
74
Double bond
Bond that forms when two electrons are shared between atoms
75
Triple bond
Bond that forms when three electrons are shared between atoms
76
Single bond
Bond that forms when just one electron is shared between atoms
77
Bonding pair
Pair of electrons shared between atoms
78
Bonding orbital
A molecular orbital that is lower in energy than any atomic orbitals from which it was formed
79
Bond order
((Number of electrons in bonding orbitals)-(Number of electrons in non-bonding orbitals))/2
Only stable when positive
80
Stable molecular bond
A molecule for which there are more bonding electrons than non-bonding electrons
81
Bond length
Average length of a bond between two particular atoms in a variety of compounds
82
Bond energy
The energy required to break 1 mol of the bond in a gasseous state
83
Chemical bond
Sharing or transfer of electrons to form stable electron configurations for bonding atoms
84
Molecular Orbital theory
Advance model of molecular bonding in which atomic orbitals are delocalized over the molecule as a whole, giving the molecule its own orbital
85
Molecular geometry
The geometric arrangement of atoms in a molecule
86
Valence shell electron pair repulsion (VSEPR) theory
Bonds themselves repel one another within a molecule
87
Electron groups
Any form of bond or even a single atom
88
Linear molecular geometry
Bonds are set up 180 degrees from one another
89
Trigonal planar geometry
Bonds are arranged 120 degrees from one another on a single plane.
90
Tertrahedral geometry
Bonds form a 4 sided pyramid shape within the molecule, bonds are spaced 109.5 degrees from one another
91
Bipyramidal geometry
Within the molecule, 3 bonds exist on one plane 120 degrees from one another, with an additional bond at either end spaced 90 degrees from the plane
92
Octahedral Geometry
In a molecule, four of the bonds exist on a plane while another rests at either end of and 90 degrees from this plane
93
Electron geometry
Geometrical arrangement of the electron groups
94
Seesaw effect
In what would have been a triangular bipyramid structure, the remaining lone pair of electrons exist as a 'bond', forming a see-saw shape
See book
95
T-shape effect
The tendency for atomic bonds to form a T-shape, when left with two electron pairs, full geometric shape would be the 'triangle bipyramidal geometry'
96
Square pyramidal
Molecular geometry of what would have been an ocahedron, with the lone pair of electrons at the bottom
97
Square planar
Molecular geometry of what would have been an ocahedron, with one lone pair of electrons at the bottom and anther lone pair of electrons at the top
98
Determining the geometry of complex molecules
1) List all of the elements present in the molecule
2) list all of the electron groups for each element type
3) list how many of those electron groups are lone pairs
4) record the best molecular geometry for each element type
99
Bent
Molecular geometry resulting from a molecule with two lone pairs and two electron groups
100
Polarity
If the entire molecule where a sphere, would one side of the sphere have more electrons than the other? If so, that side of the molecule has an overall negative charge while the opposite side has a positive charge and the entire molecule is polar.
101
Valance bond theory
A chemical bond results when valance electrons from one atom are donated to the unfilled orbitals of another atom
102
Hybridization
Mathmatical proceedure in which the individual orbitals of an atom come together to form new orbitals for the whole atom (hybrid orbitals)
1) the number of standard atomic orbitals added together always equals the number of hybrid orbitals formed.
2) the partial combonation of standard atomic orbitals added together determines the shapes and energies of the hybrid orbitals formed
3) the particular type of hybridization that occurs is the one that yields the lowest overall energy for the molecule
103
Hybrid orbitals
New orbitals made from the combination of atomic orbitals when an electron is shared between them
104
sp^3 hybridization
When a 2s and all three 2p sub-shells come together to form four sp^3 subshells, in tetrahedral geometry with 4 electron groups
105
sp^2 hybridization
A 2s and 2 of the 2p-orbitals come together in triangular planar geometry to form 3 sp^2 orbitals, leaving one p-subshell with 3 electron groups
106
Sp hybridization
When a 2s and a 2p forms 2 sp orbitals and leaves 2 p-orbitals in linear geometry, with 2 electron groups
107
Sp^3d hybridization
When a 3s, all three 3p, and a 3d form 5 new sp^3d hybrid orbitals, leaving four 3d orbitals, in triangular bipyramidal geometryo
108
Sp^3d^2 hybridization
When a 3s, all three 3p, and two 3d form 6 new sp^3d^2 hybrid orbitals, leaving thee 3d orbitals, in octahedral geometry
109
Structural isomers
Molecules with the same molecular formula but different structures
110
Dipole moments
A measure of the separation of positive and negative poles in a molecule
111
Polyatomic molecules
Molecules composed of three or more elements
112
Radio activity
The emmision of subatomic particles or high energy electromagnetic radiation by the nuclei of an atom
113
Radioactive
A term reserved for atoms whose nuclei emit subatomic particles
114
Phosphorescence
Long-lived emmision of light following the absorption of light
Glow-in-the-dark-effect
115
Types of radio activity
```
Alpha decay
Beta decay
Gamma decay
Ray emmition
Positron emmition
```
116
Nuclide
A isotope of a particular variety in terms of an element
117
Alpha decay
An unstable nucleus emmits particles composed of two protons and two neutrons (He-4)
118
Nuclear equation
Way of representing alpha decay
| (Parent isotope) → (daughter particle)+'He-4'
119
Ionizing power
The ability of radiation to ionize other molecules, alpha radiation is highest
120
Penetrating power
Opposite of ionizing power
Ability for radiation to penetrate matter
Alpha radiation particles (He-4) are too big to penetrate matter
121
Beta decay
An unstable nucleus that emmits electrons
122
Beta particle
An electron described as isotope notation
0
-1 e
123
Beta radiation equation
(Old isotope)→(new isotope)+e
124
Gamma ray emission
Electromagnetic frequency in the gamma range
Symbolized
0
0 γ
125
Positron emmition
An unstable nucleus emits a positron
126
Positron
Anti-particle of the electron
0
+1 e
127
Electron capture
When an nucleus assimilates an electron from an inner orbital of its electron cloud
128
Strong force
Natural force that binds the nucleus of an atom together
129
Nucleons
Protons and neutrons that make up a nucleus
130
N/Z ratio
Ratio of neutrons to protons
131
High N/Z ratio
Nucliedes lie above the valley of stability
132
Low N/Z
Nucliedes lie below the valley of stability
133
Nucleides
Radioactive isotopes
134
Valley of Stability
Proper ratio of neutrons to protons for stability, individual for each element
135
Magic number atoms
Atoms that contain
2, 8, 20, 28, 50, 82 neutrons
126 protons
Are always uniquely stable
136
Film-badge dosimeters
Photographic film encased in plastic and pinned to clothing
| Simple radiation detector
137
Geiger-muller counter
Particles pass through an argon chamber creating a trail of ionized argon atoms
138
Scintillation counter
Emmitions pass through a material that emmits light in the presence of radio activity
139
Rate of decay
Rate=k*N
K is the decay rate
140
Half-life of decay
T sub-1/2= .693/k
k is the decay rate
141
Integrated rate law
Ln Nt/N0 = -k*t
Nt is the number of radioactive nuclei at a time
N0 is the initial number of radioactive nuclei
142
Radiocarbon dating
Devised by Willard Libby in 1949
Used to estimate the ages of fossils and artifacts
Convert halflife to 'k'-rate
Use integrated rate law
143
Uranium/Lead Dating
Dating method used for non-living objects or objects older than 50,000 years
144
Nuclear fision
The splitting of the uranium atom
145
Chain reaction
Neutrons produced by the fission of one uranium nucleus induces the fission in other uranium nuclei
146
Critical mass
Enough U-235 to produce a self sustaining reaction
147
Nuclear reactor
Method of electricity production powered by nuclear fision to produce steam
148
Converting energy to mass
E=mc^2
149
Mass defect
The phenomenon in which an isotope-particle has a mass less than the sum of their individual components
150
Nuclear binding energy
The energy required to break a nucleum into individual nuclei
151
Binding energy curve
Graphical relationship that describes how mass relates to the binding energy per nucleon
152
Nuclear Fusion
The combonation of two lighter nuclei to form a heavier nuclei
153
Transmutation
The transformation of one element into another
154
Linear Accelerator
A device in which a charged particle is accelerated
155
Cycletron
A charged particle is accelerated back and forth between two chambers of accelerated voltage
156
Positron emission tomography
The use of positron emitting nucleides as an imaging technique
157
Radiotracer
Diagnosis technique in which a radioactive nuclide attatched to a compound or introduced into a mixture in order to track the movement of the diagnosis or mixture within the body.
158
Radiotherapy
The used of radiation to kill rapidly dividing cells
159
Other uses for radioactivity
1) Kill micro-organisms
2) used to kill bacteria within food
3) used to control the population of harmful insects
160
Polyprotic Acid
Acids that contain more than one ionizable proton
161
Diprotic acid
Acid that has both a strong ionizable proton (H+) and a weak second ionizable proton
162
Particle
A singular unit of matter, whether it be a molecule, atom, ion, nucleon, or electron
163
Predicting molecular geometry
See you tube videos
164
'Coboltus' prefix
Co+2
165
'Ferrous' prefix
Fe+2
166
'Plumbous' prefix
Pb 2+
167
'Manganous' prefix
Mn 2+
168
'Mercuric' prefix
Hg 2+
169
'Nickelous' prefix
Ni 2+
170
'Stannous' prefix
Sn 2+
171
'Cobaltic' prefix
Co 3+
172
'Ferric' prefix
Fe 3+
173
'Manganic' prefix
Mn 3+
174
'Plumbic' prefix
Pb 4+
175
'Stannic' prefix
Sn 4+
176
'Amonium' prefix
NH4
177
'Hydronium' prefix
H3O+
178
'Mercurous' prefix
Hg2 2+
179
'Acetate' suffix
C2H3O2-
180
'Bicarbonate' suffix
HCO3-
181
'Bisulfate' suffix
HSO4-
182
'Chlorate' suffix
ClO3-
183
'Chlorite' suffix
ClO2-
184
'Dihydrogen phosphate' suffix
H2PO4-
185
'Hydroxide' suffix
OH-
186
'Hypochlorite' suffix
ClO-
187
'Nitrate' suffix
NO3-
188
Metalic -suffix
-ide
189
'Nitrite' suffix
NO2-
190
'Perchlorate' suffix
ClO4-
191
'Permanganate' suffix
MnO4-
192
'Biphosphate' suffix
HPO4 2-
193
'Carbonate' suffix
CO3 2-
194
'Chromate' suffix
CrO4 2-
195
'Dichromate' suffix
Cr2O7 2-
196
'Peroxide' suffix
O2 2-
197
'Sulfate' suffix
SO4 2-
198
'Sulfite' suffix
SO3 2-
199
'Thiosulfate' suffix
S2O3 2-
200
'Phosphate' suffix
PO4 3-
201
'Phosphite' suffix
PO3 3-
202
Binary (H+)(_ -) acid
'Hydro_ic acid'
203
Oxy acid nomenclature (the anion ends in oxygen)
| H+)(_O
- Ate suffix....'_ic acid'
| - ite suffix...'_ous acid'
204
Cone-shaped bond
Look it up
205
Dashed bond
Look it up
206
Partial Pressure
P.solvent=(n-fraction.solvent)*P.solvent
| P.solute=(n-fraction.solute)*P.solute
207
Friability
The ability of a substance to be easily broken down into smaller pieces
