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Flashcards in Structure Of Matter Deck (207):
0

Antoine Lavoisier

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

2

Law of conservation of mass

In a chemical reaction, matter is niether created nor destroyed
Key concept to the formulation of modern atomic theory

3

Joseph Proust

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

4

Law of definite proportions

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

5

Mass ratio

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

John Dalton

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

Law of multiple proportions

When two elements form two different compounds the masses of the elements can be described in a mass ratio

8

Atomic theory

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

Atomic mass

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

Physical determination of atomic mass

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

11

Mass spectrometery

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

Atomic number (z)

Tells
1) number of protons
2) number of electrons (when nuetral)
3) positive charge of the nucleus
Calculated as z=A-N

13

Neutron Count (N)

Number of neutrons in a given atom
Calculated as N=A-z

14

Mass number (A)

Total number of nucleons (protons and neutrons) in an atom
A=z+N
Often written as
[Element name]-A
(eg: carbon-12)

15

Isotopes

Atoms with the same number of protons but different neutrons

16

Isotope notation

A over z and then the element name
Mass number is always written over the number of protons

17

Natural abundance

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

18

Electron energy levels

Each row of the periodic table is a new energy level for electrons
The 'unfilled level' being the level containing valance electrons

19

Electron configuration

Way of showing which electrons occupy which orbitals
Ground state= H 1s^1

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Pauli exclusion principal

No two electrons can ever have the same four quantum numbers

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Orbital

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

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P orbital

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

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D orbital

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

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F orbital

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

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Quantum numbers

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

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S orbital

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

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Principal quantum number (n)

Positive integer that represents electron energy level

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Angular momentum quantum number(l)

Integer that represents the shape of the orbital

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

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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

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Dmitry Mendeleev

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

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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

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Aufbau Principal

Only two electrons of opposite spins are allowed in each orbital

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Hunds Rule

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

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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

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Covalent radius

Radius of an atom when it is bonded to another atom

51

Atomic radii

Lower left of the periodic table

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Electronegativity

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

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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

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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

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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

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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

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Lewis theory

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

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Lewis structures

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

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Octet rule

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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Lewis base

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

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Lewis acid

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

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Double bond

Bond that forms when two electrons are shared between atoms

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Triple bond

Bond that forms when three electrons are shared between atoms

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Single bond

Bond that forms when just one electron is shared between atoms

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Bonding pair

Pair of electrons shared between atoms

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Bonding orbital

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

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Bond order

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

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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

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Bond energy

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

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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

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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+

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'Plumbic' prefix

Pb 4+

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'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