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Flashcards in States Of Matter Deck (141):
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Pressure

The distribution of a force across a designated area
P=Fn/A

1

MmHg

Unit of measure for pressure based on barometer readings

2

Torr

Unit of pressure equal to 1mmHg

3

Pascal (Pa)

Unit of pressure
Pa= 1 N/m^2

4

atm

Unit of pressure measure based on sea level
Generally accepted
=101,325 Pa
=29.92 mmHg
=14.7 psi

5

Manometer

Laboratory equipment used to measure pressure

6

Simple gas laws

1) boyles law
2) charles law
3) avogadros law

7

Boyles law

V1*P1=V2*P2

8

Charles' law

V1/T1=V2/T2

9

Avagodro's law

V1/n1=V2/n2

10

Ideal gas law

P*V=n*R*T

R=0.08206

atm, L, mol, K

11

Ideal Gas Constant

R=0.08206

12

Molar Volume

Volume occupied by one mole of a substance

13

Standard temperature

0 C = 273 K

14

Standard pressure

1 atm

15

Gas density

Density = (molar mass)/(molar volue)

16

Partial pressure

Presure exerted by an individual gas in a mixture

Pn=nn (RT/V)

17

Daltons law

The sum of all partial pressures is equal to the total pressure

18

Kinetic molecular theory

Gas is represented by a series of tiny particles
1) the size of a particle is negligably small
2) the kinetic energy of each particle is proportional to the temperature in Kelvins
3) The collision of one particle with another results in an exchange of energy

19

Work on the Kinetic molecular theory with the gas law

226

20

Avagadros number

NA= 6.022x10^-23

21

Mole

Avagodros number of atoms (or molecules) of a given substance

22

Kinetic energy of a given particle

KE=1/2 (m*v^2)

Lighter particles carry more kinetic energy

23

Average kinetic energy

KE=(3/2)*R*T

24

Mean free path

The average distance a particle travels between collisions

25

Diffusion

The process by which gas molecules spread out in response to a concentration gredient

26

Effusion

Process by which gas escapes from a container into a vacuum through a small hole

27

Graham's law of effusion

Describes the relationship between effusion rates

RateA/RateB = sqrt(molar massB/molar massA)

28

Gas law correction for volume

V=nRT/P+ nb
b is a constant represented by individual gasses

29

Gas law correction for pressure

P=nRT/V- a(n/V)^2
a is a constant represented by individual gasses

30

Vanderwals equation

[P+a(n/V)^2]x[V-nb]=n*R*T

Product of Corrected gas volume and corrected Volume equals the rest of the ideal gas law

31

Pressure-Volume relationship to atom count

PV=NkT

N=number of atoms
k=boltzman constant=1.38x10^-23

32

Boltzman constant

1.38x10^-23

33

Liquid

Definite volume
Assumes the shape of a container

34

Solid

Definite Volume
Definite Shape
High density
Particles packed closely together
Low kinetic energy

35

Gas

No definite volume,
No definite shape
Low density
Maximum space between particles
High kinetic energy

36

Melting

Physical change from solid to liquid

37

Evaporation

Physical change from liquid to gas

38

Sublimation

Physical change from solid directly to gas

39

Condensation

Physical change from gas to liquid

40

Freezing

Physical change from liquid to solid

41

Deposition

Physical change from gas to solid

42

Intermolecular forces

Forces that hold condensed (non-gaseous) states together

43

Dispersion force

Result of fluctuations in the electron distribution within molecules and atoms

44

Dipole-dipole forces

In polar molecules, forces that keep maintain the molecular geometry of a molecule

45

Miscilibility

The ability to mix without separating into separate states

46

Ion-dipole forces

Result of a polar compound mixing with an ionic compound

47

Surface tension

Energy required to increase the area of the liquid by a given unit

48

Viscocity

The resistance of a liquid to flow

49

Capillary action

The ability of a liquid o flow against gravity up a narrow tube

50

Volatile

A liquid's ability to vaporize easily

51

Nonvolatile

Refers to liquids that do not vaporize easily

52

Phase Diagram

Pressure/temperature chart in which the phases are mapped out with the fusion curve, sublimation curve, and vaporization curve with the triple point where they all meet

53

Critical point

The point of maximum temperature and pressure at which substances become a superficial liquid, neither liquid nor gas

54

Sublimation curve

The line on the phase diagram lowest on the pressure/temperature chart
Higher pressure yeilds solid
Higher temperature yeilds gas

55

Fusion curve

The line that separates liquids from solids on the phase diagram
Runs up from the 'triple point'
High temperature yields liquid
Low temperature yields a solid

56

Vaporization curve

The line on the phase diagram that seperates liquid from gas
Splits right from the triple point
High pressure yields liquid
High temperature yields gas

57

Triple point

The point at which the temperature pressure ratio is at a delacate balance, where the smallest change in either variable could change the phase of the substance

58

Crystalline structure

A solid in which atoms molecules or ions are arranged in patterns with long range repeating order

59

Solution

A homogeneous mixture of two substances

60

Solvent

Majority component in a solution

61

Solute

Minority component of a solution

62

Aqueous Solutions

A solution in which water is the solvent

63

Solubilty

The amount of a substance that will dissolve in a given amount of solvent

64

Natures mixing tendency

The resulting solution will be the lowest energy level possible

65

Entropy

A measure of energy dispersal within a system

66

Properties of the ideal gas

Molar Volume 22.41L at STP
No adjustment for VP=nRT
(V/nRT)=1

67

Miscible substances

Substances that are soluble in one another

68

Exothermic

Heat is given off, feels hot

69

Endothermic

Absorbs thermal energy, feels cold

70

Solution prosses

1) Solute is broken into constituent particles (endothermic)
2) The solvent particles separate to make room for the solute particles (endothermic)
3) Solute are dispersed throughout solvent (exothermic)

71

Enthalpy of solution

The sum of the enthalpies for each step in the solution process
ΔH=Δh-solute seperation+Δh-solvent seperation+Δh-mixture

72

Heat of hydration

The sum of the heat of solvent separation (endothermic) and and the heat of the mixture (exothermic)

73

Dynamic equilibrium

The point at which the rates of dissolution and recrystaization become equal

74

Saturated solution

Solution in which the amount of dissolvable solute has already been reached,
If any more solute is added, it will not dissolve

75

Nonsaturated solution

A solution that could still dissolve more solute

76

Supersaturated solution

Solution containing more solute than could be disolved in that solvent

77

Recrystalization

A purification technique by which a solid is added to water to create a saturated solution at a high temperature, as the solution cools, it becomes over saturated and produces crystals of the dissolved solid

78

Factors in solubility

1) Volume
2) Temperature
3) Pressure

79

Henrys law

Solubility of a gas (units of M) is equal to the product of henrys law constant (k sub-h, individual to each gas) and the partial pressure of that particular gas

S=k*P

80

Dilute solution

Solution containing some small quality of solute relative to the amount of solvent

81

Concentrated Solution

A solution containing a large quantity of solute, relative to the solvent

82

Molarity (M)

(Moles of solute)/(volume of the solution in L)

83

Molality (m)

(Moles of solute)/(mass of solvent in Kg)

84

Percent by mass

Way of representing the saturation of a solution

(Solute mass)/(solution mass)x100%

85

Parts per million (ppm)

(Mass solute/mass solution)x10^6

86

Parts per billion (ppb)

(Mass solute/mass solution)x10^9

87

Parts by volume

(Volume solute)/(volume solution)x(multiplication factor)

88

Usage of parts by mass or volume

Find the mass or volume of the solution given the solute specifications (vice versa)

89

Mole fraction

(Moles Solute)/(moles of solute and solvent added) x100%

90

Colligative property

Any property that depends on the number of particles dissolved in a solution

91

Raoult's law

The vapor pressure of the entire solution is equal to the product of the mole fraction and the vapor pressure of the pure solvent

See page 571 for practice

92

Ideal solution

When both the solute and solvent are volitile, contributing to the vapor pressure of the overall solution

93

Freezing point depression

lowering the freezing point of a solution,
Colligative property

ΔTf=m*Kf
m-molality
Kf freezing point depression constant for the solvent

94

Boiling point elevation

Ease of Raising the boiling point of a solution,
Colligative property

95

Osmosis

The flow of a solvent from a solution of lower solute concentration

96

Semipermeable membrane

A material that effectively filters some substances but not others

97

Osmotic pressure

Equal to molarity times the gas constant times solution temperature

M*R*T

98

Van't Hoff factor

i=(moles of particles in solution)/(moles of formula units dissolved)

99

Colloids

A substance in which fine particles (between 1nm 1000nm in size) are evenly dispersed,
ie Milk

100

Micelles

Non-polar hydrocarbon trails crowd into the center of a shape to maximize their interactions with one another

101

Tyndall effect

The scattering of light by colloidal dispersion

102

Moles present

(mass present)/(substance molar-mass)

103

Aqueous solution

A solution in which water acts as the solvent

104

Dilution equation

M1*V1=M2*V2

105

Electrolytes

Solutes that dissolve into ions that conduct electricity

106

Strong electrolytes

Electrolytes that dissolve completely in water

107

Non-electrolytes

Solutes whose ions do not conduct electricity

108

Strong acid

An acid that completely ionizes in water

109

Weak acids

An acid that does not completely ionize in water

110

Weak Electrolytes

Electrolytes that that do not dissolve completely in water

111

Solubility rules

1) Group1 and Amonia cations
2) Nitrates and acetates
3) Chlorides, Bromides, Iodines (Ag1, Hg2, Pb2)
4) Sulfate (Sr2, Ba2, Pb2, Ag1, Ca2)
5) But not hydroxide or sulfide (Unless rule 1 applies, Ca2, Sr2, or Ba2)
6) But not carbonates or phosphates (unless rule 1)

112

Precipitation reactions

When a solid (precipitate) is formed by the mixing of two solutions

113

Precipitate

Solid formed by mixing two solutions

114

Acid solutions

Acids release H+ ions in the presence of water

115

Basic Solutions

Bases release OH- in the presence of water

116

Hydronium ions

H+(aq) ions released by an acid in water

117

Titration

Substance in a solution of a known concentration is reacted with another substance in a solution of unknown concentration

118

Equivilence point

In an acid base reaction, the Point at which moles of OH- equal moles of H+

119

Indicator

A dye used that changes color based on acidic exposure, used to indicate equivalence point of a solution

120

Structural Isomers

Atoms are connected to one another in a variety of ways

121

Stereoisomers

Atoms are connected in the same way, with a variety of specially arrangements about the metal atom

122

Coordination isomers

Occurance of a coordinated ligand exchanging places with an uncoordinated counterion

123

Linkage isomers

Stereoisomers that have ligands that coordinate to the metal in different geometric orientations.

124

Geometric isomers

A type of stereoisomerism that results from ligands bonding to metal in a variety of spacial arrangements

125

Hydroscopy

The ability of a substance to pull moisture from the surrounding environment

126

Crystallgraphic deffect

Interruptions in the regular pattern of fixed repeating order solids (crystals)

127

Microcrystaline

A solid that contains crystals visible only through microscopic examination

128

Cubic crystal system

A crystal system in which the unit cell is in the shape of a cube
Toms are arranged in cubic order

129

Crystal family (lattice system)

Several classes of crystaline structures that have similar symmetries

130

Is it polar (from the chemical formula)

Determine the molecular geometry
Is it has lone pair electrons, it is polar

131

Cone shaped bond in complex molecules

Look it up

132

Dashed bond in complex molecules

Look it up

133

Complex molecular orbit calculations

Practice

134

Disiccant

Hygroscopic material that induces dryness in surroundings

135

Brauner-Emmet-Teller (BET) theory

Solids are capable of absorbing gas particles at their surface

136

Langmuer Equation

Reated the quantity of gas molecules absorbed by the surface of a solid to the gas pressure

137

Calculating Ph value from hydrogen ion activity

Ph=log(1/[H+])
Where h+ refers to the hydrogen ion activity of a solution

138

Supercritical drying

Transitioning a material from a liquid to a gas by raising the temperature and pressure beyond the supercritical point, and then lowing the pressure until the substance become a gas

139

Surface energy

The quantitive measurement of the bonds disrupted by the formation of a surface

140

Excimer

'Excited dimer'
A short-lived molecule, formed from two species, with a filled vallance shell (noble gas)