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A) Chemistry > Structure Of Matter > Flashcards

Structure Of Matter Flashcards

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1
Q

Antoine Lavoisier

A

Credited with the formulation of the law of conservation of matter in 1789

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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

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3
Q

Joseph Proust

A

Credited with the formulation of the law of definite proportions in 1797

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4
Q

Law of definite proportions

A

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

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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

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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)

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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

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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

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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

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10
Q

Physical determination of atomic mass

A

Sum of the products of each isotope’s decimal abundance, mass of each isotope

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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

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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
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13
Q

Neutron Count (N)

A

Number of neutrons in a given atom

Calculated as N=A-z

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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)
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15
Q

Isotopes

A

Atoms with the same number of protons but different neutrons

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16
Q

Isotope notation

A

A over z and then the element name

Mass number is always written over the number of protons

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17
Q

Natural abundance

A

Percentage breakdown of isotopes present in any particular sample of an element, varies by isotope

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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

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19
Q

Electron configuration

A

Way of showing which electrons occupy which orbitals

Ground state= H 1s^1

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20
Q

Pauli exclusion principal

A

No two electrons can ever have the same four quantum numbers

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21
Q

Orbital

A

Any one of four areas that describe the entire distribution of where that electron might be located.

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22
Q

P orbital

A

Electrons here occupy an any space within a peanut shape arround the nucleus

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23
Q

D orbital

A

Electrons here occupy an any space within a flower shape arround the nucleus

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24
Q

F orbital

A

Electrons here occupy an any space within a 3D-flower shape arround the nucleus

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25
Q

Quantum numbers

A

Values that describe the behavior of an individual electron within an atom

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26
Q

S orbital

A

Lowest energy

Electrons here may occupy a sphere-like area around the nucleus

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27
Q

Principal quantum number (n)

A

Positive integer that represents electron energy level

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28
Q

Angular momentum quantum number(l)

A

Integer that represents the shape of the orbital

S=0
P=1
D=2
F=3

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29
Q

Magnetic quantum number (m sub-l)

A

Integer that represents orbital orientation
Exist on the interval (-l to +l)
l is the angular momentum quantum number

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30
Q

Spin quantum number (m sub-s)

A

Refers to the direction in which the electron spins

Either -1/2 or +1/2

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31
Q

Periodic property

A

A property of an atom predictable by its position on the periodic table

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32
Q

Electron configuration

A

A way of writing which electrons are in which orbital

‘Ground State’- 1s^1

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33
Q

Dmitry Mendeleev

A

Russian chemist Credited with the arrangement of the periodic table based on the work of German chemist Julius Lothar Meyer

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34
Q

Orbital Diagram

A

Way of drawing the ‘direction’ of electron ‘spin’ within electron configuration
Little arrows going up and down

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35
Q

Pauli Exclusion Principal

A

No two electrons can have the same four quantum numbers within the same atom

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36
Q

Coulombs law

A

E=q1q2/r1/(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

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37
Q

Shielding

A

The property of an electron to repel another electron

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38
Q

Effective Nuclear charge

A

Refers to the charge acting on a particular electron

The ionic charge of that atom if the electron did not exist

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39
Q

Penetration

A

The action of an electron changing energy levels due to an attraction to the Nucleus

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40
Q

Aufbau Principal

A

Only two electrons of opposite spins are allowed in each orbital

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41
Q

Hunds Rule

A

Electrons must fill these orbitals as single electrons before they can double up

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42
Q

Valance Electrons

A

Electrons that are exchanged during chemical bonding

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43
Q

Core electrons

A

Electrons that would compose an atom of the noble nearest two but less than the mass of the atom in question

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44
Q

Orbital Blocks of the periodic table

A

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

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45
Q

Noble gasses

A

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

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46
Q

Alkali Metals

A

Group 1 elements, highly reactive
One valence electron
Excelent reducing agents (decrease charge of a material by contributing an electron)
Result in violent reactions

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47
Q

Alkali Earth Metals

A

Group 2 elements

Two valence electrons

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48
Q

Halogens

A

Group 7
7 valence electrons
Excelent oxideizing agents (take electrons from a material thereby increasing their charge)

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49
Q

Van der wals radius

A

The radius of an atom when it is not bonded to another atom

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50
Q

Covalent radius

A

Radius of an atom when it is bonded to another atom

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51
Q

Atomic radii

A

Lower left of the periodic table

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52
Q

Electronegativity

A

Ability for an atom to attract electrons
Opposite of metallic Character
Increases to the upper right of an atom

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53
Q

Metallic Character

A

Ability for an atom to loose an electron
Opposite of electronegativity
Increases to the lower left of the periodic table

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54
Q

Paramagnetism

A

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

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55
Q

Diamagnetic

A

Property of atoms to be repelled by an external magnetic field
An effect of full energy orbitals
Usually attributed to zinc

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56
Q

Cation radius rule

A

Cations are nearly half the size of their nuetral atoms

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57
Q

Anion radius rule

A

Anions are much larger than their neutral atoms

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58
Q

Electron affinity

A

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

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59
Q

Ionic bonds

A

Oppositely charges ions attract one another for an exceptionally low net charge
Electrons are tranfered in this bond
Electronegativity <.5

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60
Q

Covalent bonds

A

Two nonmetals bonding together
Form molecules
Atoms share their electrons
Electronegativity .5

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61
Q

Metallic bonds

A

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

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62
Q

Macromolecular bonding

A

Smaller molecules bonded together in a polymer chain

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63
Q

Hydrogen bonding

A

Strong dipole-dipole bond between Hydrogen and either O, N, or F

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64
Q

Polar covalent bonds

A

High electronegativity difference between two atoms bonded together
Electronegativity > 1.7

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65
Q

Pi bond

A

Bond that forms betwwen two overlapping P orbitals

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66
Q

Peptide bond

A

Bond that forms between the amine end of one amino acid and the carboxylic end of another

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67
Q

Sigma bond

A

Bond resulting from an overlap of a p orbits and an s orbital
Commonly sp^2 hybridization

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68
Q

Valence bond theory

A

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

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69
Q

Lewis theory

A

Simple model of of chemical bonding in which atoms are arranged to form ‘octets’ for each atom

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70
Q

Lewis structures

A

Drawings of atoms in which bonds are represented as dots to form octets

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71
Q

Octet rule

A

When bonded together, electrons must share as many electrons as it takes so that they each have ready access to 8 electrons

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72
Q

Lewis base

A

An atom, ion, or molecule that donates an electron pair

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73
Q

Lewis acid

A

An atom, ion, or molecule that accepts an electron pair

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74
Q

Double bond

A

Bond that forms when two electrons are shared between atoms

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75
Q

Triple bond

A

Bond that forms when three electrons are shared between atoms

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76
Q

Single bond

A

Bond that forms when just one electron is shared between atoms

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77
Q

Bonding pair

A

Pair of electrons shared between atoms

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78
Q

Bonding orbital

A

A molecular orbital that is lower in energy than any atomic orbitals from which it was formed

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79
Q

Bond order

A

((Number of electrons in bonding orbitals)-(Number of electrons in non-bonding orbitals))/2
Only stable when positive

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80
Q

Stable molecular bond

A

A molecule for which there are more bonding electrons than non-bonding electrons

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81
Q

Bond length

A

Average length of a bond between two particular atoms in a variety of compounds

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82
Q

Bond energy

A

The energy required to break 1 mol of the bond in a gasseous state

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83
Q

Chemical bond

A

Sharing or transfer of electrons to form stable electron configurations for bonding atoms

84
Q

Molecular Orbital theory

A

Advance model of molecular bonding in which atomic orbitals are delocalized over the molecule as a whole, giving the molecule its own orbital

85
Q

Molecular geometry

A

The geometric arrangement of atoms in a molecule

86
Q

Valence shell electron pair repulsion (VSEPR) theory

A

Bonds themselves repel one another within a molecule

87
Q

Electron groups

A

Any form of bond or even a single atom

88
Q

Linear molecular geometry

A

Bonds are set up 180 degrees from one another

89
Q

Trigonal planar geometry

A

Bonds are arranged 120 degrees from one another on a single plane.

90
Q

Tertrahedral geometry

A

Bonds form a 4 sided pyramid shape within the molecule, bonds are spaced 109.5 degrees from one another

91
Q

Bipyramidal geometry

A

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
Q

Octahedral Geometry

A

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
Q

Electron geometry

A

Geometrical arrangement of the electron groups

94
Q

Seesaw effect

A

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
Q

T-shape effect

A

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
Q

Square pyramidal

A

Molecular geometry of what would have been an ocahedron, with the lone pair of electrons at the bottom

97
Q

Square planar

A

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
Q

Determining the geometry of complex molecules

A

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
Q

Bent

A

Molecular geometry resulting from a molecule with two lone pairs and two electron groups

100
Q

Polarity

A

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
Q

Valance bond theory

A

A chemical bond results when valance electrons from one atom are donated to the unfilled orbitals of another atom

102
Q

Hybridization

A

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
Q

Hybrid orbitals

A

New orbitals made from the combination of atomic orbitals when an electron is shared between them

104
Q

sp^3 hybridization

A

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
Q

sp^2 hybridization

A

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
Q

Sp hybridization

A

When a 2s and a 2p forms 2 sp orbitals and leaves 2 p-orbitals in linear geometry, with 2 electron groups

107
Q

Sp^3d hybridization

A

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
Q

Sp^3d^2 hybridization

A

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
Q

Structural isomers

A

Molecules with the same molecular formula but different structures

110
Q

Dipole moments

A

A measure of the separation of positive and negative poles in a molecule

111
Q

Polyatomic molecules

A

Molecules composed of three or more elements

112
Q

Radio activity

A

The emmision of subatomic particles or high energy electromagnetic radiation by the nuclei of an atom

113
Q

Radioactive

A

A term reserved for atoms whose nuclei emit subatomic particles

114
Q

Phosphorescence

A

Long-lived emmision of light following the absorption of light
Glow-in-the-dark-effect

115
Q

Types of radio activity

A 
Alpha decay
Beta decay
Gamma decay
Ray emmition
Positron emmition
116
Q

Nuclide

A

A isotope of a particular variety in terms of an element

117
Q

Alpha decay

A

An unstable nucleus emmits particles composed of two protons and two neutrons (He-4)

118
Q

Nuclear equation

A

Way of representing alpha decay

(Parent isotope) → (daughter particle)+’He-4’

119
Q

Ionizing power

A

The ability of radiation to ionize other molecules, alpha radiation is highest

120
Q

Penetrating power

A

Opposite of ionizing power
Ability for radiation to penetrate matter
Alpha radiation particles (He-4) are too big to penetrate matter

121
Q

Beta decay

A

An unstable nucleus that emmits electrons

122
Q

Beta particle

A

An electron described as isotope notation

0
-1 e

123
Q

Beta radiation equation

A

(Old isotope)→(new isotope)+e

124
Q

Gamma ray emission

A

Electromagnetic frequency in the gamma range

Symbolized

0
0 γ

125
Q

Positron emmition

A

An unstable nucleus emits a positron

126
Q

Positron

A

Anti-particle of the electron

0
+1 e

127
Q

Electron capture

A

When an nucleus assimilates an electron from an inner orbital of its electron cloud

128
Q

Strong force

A

Natural force that binds the nucleus of an atom together

129
Q

Nucleons

A

Protons and neutrons that make up a nucleus

130
Q

N/Z ratio

A

Ratio of neutrons to protons

131
Q

High N/Z ratio

A

Nucliedes lie above the valley of stability

132
Q

Low N/Z

A

Nucliedes lie below the valley of stability

133
Q

Nucleides

A

Radioactive isotopes

134
Q

Valley of Stability

A

Proper ratio of neutrons to protons for stability, individual for each element

135
Q

Magic number atoms

A

Atoms that contain
2, 8, 20, 28, 50, 82 neutrons
126 protons
Are always uniquely stable

136
Q

Film-badge dosimeters

A

Photographic film encased in plastic and pinned to clothing

Simple radiation detector

137
Q

Geiger-muller counter

A

Particles pass through an argon chamber creating a trail of ionized argon atoms

138
Q

Scintillation counter

A

Emmitions pass through a material that emmits light in the presence of radio activity

139
Q

Rate of decay

A

Rate=k*N

K is the decay rate

140
Q

Half-life of decay

A

T sub-1/2= .693/k

k is the decay rate

141
Q

Integrated rate law

A

Ln Nt/N0 = -k*t

Nt is the number of radioactive nuclei at a time
N0 is the initial number of radioactive nuclei

142
Q

Radiocarbon dating

A

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
Q

Uranium/Lead Dating

A

Dating method used for non-living objects or objects older than 50,000 years

144
Q

Nuclear fision

A

The splitting of the uranium atom

145
Q

Chain reaction

A

Neutrons produced by the fission of one uranium nucleus induces the fission in other uranium nuclei

146
Q

Critical mass

A

Enough U-235 to produce a self sustaining reaction

147
Q

Nuclear reactor

A

Method of electricity production powered by nuclear fision to produce steam

148
Q

Converting energy to mass

A

E=mc^2

149
Q

Mass defect

A

The phenomenon in which an isotope-particle has a mass less than the sum of their individual components

150
Q

Nuclear binding energy

A

The energy required to break a nucleum into individual nuclei

151
Q

Binding energy curve

A

Graphical relationship that describes how mass relates to the binding energy per nucleon

152
Q

Nuclear Fusion

A

The combonation of two lighter nuclei to form a heavier nuclei

153
Q

Transmutation

A

The transformation of one element into another

154
Q

Linear Accelerator

A

A device in which a charged particle is accelerated

155
Q

Cycletron

A

A charged particle is accelerated back and forth between two chambers of accelerated voltage

156
Q

Positron emission tomography

A

The use of positron emitting nucleides as an imaging technique

157
Q

Radiotracer

A

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
Q

Radiotherapy

A

The used of radiation to kill rapidly dividing cells

159
Q

Other uses for radioactivity

A

1) Kill micro-organisms
2) used to kill bacteria within food
3) used to control the population of harmful insects

160
Q

Polyprotic Acid

A

Acids that contain more than one ionizable proton

161
Q

Diprotic acid

A

Acid that has both a strong ionizable proton (H+) and a weak second ionizable proton

162
Q

Particle

A

A singular unit of matter, whether it be a molecule, atom, ion, nucleon, or electron

163
Q

Predicting molecular geometry

A

See you tube videos

164
Q

‘Coboltus’ prefix

A

Co+2

165
Q

‘Ferrous’ prefix

A

Fe+2

166
Q

‘Plumbous’ prefix

A

Pb 2+

167
Q

‘Manganous’ prefix

A

Mn 2+

168
Q

‘Mercuric’ prefix

A

Hg 2+

169
Q

‘Nickelous’ prefix

A

Ni 2+

170
Q

‘Stannous’ prefix

A

Sn 2+

171
Q

‘Cobaltic’ prefix

A

Co 3+

172
Q

‘Ferric’ prefix

A

Fe 3+

173
Q

‘Manganic’ prefix

A

Mn 3+

174
Q

‘Plumbic’ prefix

A

Pb 4+

175
Q

‘Stannic’ prefix

A

Sn 4+

176
Q

‘Amonium’ prefix

A

NH4

177
Q

‘Hydronium’ prefix

A

H3O+

178
Q

‘Mercurous’ prefix

A

Hg2 2+

179
Q

‘Acetate’ suffix

A

C2H3O2-

180
Q

‘Bicarbonate’ suffix

A

HCO3-

181
Q

‘Bisulfate’ suffix

A

HSO4-

182
Q

‘Chlorate’ suffix

A

ClO3-

183
Q

‘Chlorite’ suffix

A

ClO2-

184
Q

‘Dihydrogen phosphate’ suffix

A

H2PO4-

185
Q

‘Hydroxide’ suffix

A

OH-

186
Q

‘Hypochlorite’ suffix

A

ClO-

187
Q

‘Nitrate’ suffix

A

NO3-

188
Q

Metalic -suffix

A

-ide

189
Q

‘Nitrite’ suffix

A

NO2-

190
Q

‘Perchlorate’ suffix

A

ClO4-

191
Q

‘Permanganate’ suffix

A

MnO4-

192
Q

‘Biphosphate’ suffix

A

HPO4 2-

193
Q

‘Carbonate’ suffix

A

CO3 2-

194
Q

‘Chromate’ suffix

A

CrO4 2-

195
Q

‘Dichromate’ suffix

A

Cr2O7 2-

196
Q

‘Peroxide’ suffix

A

O2 2-

197
Q

‘Sulfate’ suffix

A

SO4 2-

198
Q

‘Sulfite’ suffix

A

SO3 2-

199
Q

‘Thiosulfate’ suffix

A

S2O3 2-

200
Q

‘Phosphate’ suffix

A

PO4 3-

201
Q

‘Phosphite’ suffix

A

PO3 3-

202
Q

Binary (H+)(_ -) acid

A

‘Hydro_ic acid’

203
Q

Oxy acid nomenclature (the anion ends in oxygen)

H+)(_O

A
  • Ate suffix….’_ic acid’

- ite suffix…‘_ous acid’

204
Q

Cone-shaped bond

A

Look it up

205
Q

Dashed bond

A

Look it up

206
Q

Partial Pressure

A

P.solvent=(n-fraction.solvent)*P.solvent

P.solute=(n-fraction.solute)*P.solute

207
Q

Friability

A

The ability of a substance to be easily broken down into smaller pieces

A) Chemistry (9 decks)

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