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Chem/Phys Exam 2 > Chem Phys Exam 2 > Flashcards

Chem Phys Exam 2 Flashcards

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1
Q

Molecular interactions often occur in _____________

A

Solutions

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2
Q

Solution

A
  • Homogenous mixture in which one substance is dispersed in another as individual atoms or molecules
  • liquid solution is transparent
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3
Q

Clear Solution

A
  • When you have a solution and you have a solute it disperses to individual atoms or molecules
  • When this happens, molecules are so small they do NOT bend light, so solution is clear
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4
Q

Cloudy Liquid

A
  • Large particle that refract light
  • Particles stick together- make cloudy liquid
  • NOT considered a solution
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5
Q

Sodium Chloride

A

Ion-Ion interaction

NaCl——> Na+ Cl-

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6
Q

NaCl in solution

A
  • Ion- diploe interaction
  • Separates into anion and water molecules=partial positive
  • Proton portion=partial positive charge
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7
Q

Acid

A

Compound in which hydrogen ions are associated with an anion
Disassociates in solution:
1. Hydrogen ion (H+)
2. Anion

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8
Q

Mixtures that are NOT Solutions

A
  • Cloudy liquid (solutions are clear)
  • gasoline and water
  • oil and vinegar
  • Diprivan (propofol)
  • Maalox
  • NPH insulin (mixture and solution)
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9
Q

Gasoline and Water

A

-Most important of polar ends=do NOT mix
-oil is less dense, floats to top and water is on bottom
-Not enough interactions between these two for them to come
Together, react, and stay together

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10
Q

Concentrations and Dilutions

A
  • Molarity: best way to capture measurement
  • Molality
  • Normality
  • Equivalents
  • Osmolarity
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11
Q

Concentrated Solution

A
  • Color will give you idea about how concentrated a solution is
  • Molecules are not bound together, but they are closer to each other
  • More chemical reactions occur in concentrated solutions than in dilute solutions
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12
Q

Dilute Solution

A
  • Less chemical reactions occurring in this solution

- Molecules or atoms are farther apart

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13
Q

Mole

A
  • Amount of a substance equal to its formula weight in grams

- How we quantify all of the molecules and atoms we cannot see

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14
Q

Example of Atomic Weight: 1 Mole of atoms

A

Hydrogen= Atomic weight 1 1gram (1 mole of atoms)
Carbon= Atomic weight 12: 12 grams (1 mole of atoms)
Nitrogen=atomic weight 14 14 grams (1 mole of atoms)

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15
Q

Molecular weight of H2 & 1 mole of H2

A

Molecular weight= 2

1 mole of H2= 2 grams

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16
Q

Molecular weight of N2 & 1 mole of N2

A

Molecular weight of N2= 28

1 mole of N2= 28 grams

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17
Q

Mole

A

Weight in grams of a substance equal to its:

  1. Formula Weight
  2. Molecular Weight
  3. Gram molecular weight
  4. Atomic Mass Units
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18
Q

Avogadro’s Number

A
  • Number of atoms (for an element) or molecules (for a compound) in a mole
  • 1 mole contains 6.02 X 10^23
  • This number does NOT change
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19
Q

Molarity

A

Means of expressing the concentration of a certain species in solution in units of moles per liter

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20
Q

Molarity- How to calculate 1 Molar Solution

A
  • 1 Molar Solution (1 M)
  • Weight out solute as precisely as possible (molecular weight in grams)
  • Add enough water to make 1 liter as precisely as possible
  • want to make 1.000 M solution
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21
Q

Molarity example: NaHCO3

A
  • Molecular weight: (23+1+12+16x3)
  • 84 gram/liter=1 M
  • 84 mg/liter= 1mM - 1milimolar (1,000X less)
  • Usually talk about milimolar in chemistry
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22
Q

Calculations to Know: Molarity

A

Molarity=Moles per liter

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23
Q

Calculations to know: Moles

A

Moles= weight/ weight per mole

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24
Q

Calculations to know: Weight

A

Weight= moles X weight per mole

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25
Q

Calculations to know: Concentration

A

Concentration=moles/volume

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26
Q

Scientific Notation

A

Useful for expressing numbers that are very large and very small

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27
Q

Exponents: 10^0, 10^1, 10^2

A 
10^0= 1 (zero zeros)
10^1= 10 (1 zero)
10^2= 100 (2 zeros)
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28
Q

To determine the exponent

A
  • Count number of zeros after the 1

- Count number of places decimal point has been displaced from 1

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29
Q

Determine 10,000= 10^?

A

Count number of zeros=4

10,000= 10^4

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30
Q

Exponents of numbers

A

-exponents of numbers

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31
Q

Negative Exponents:
.1
.01
.001

A 
.1= 10 ^-1= 1/10
.01= 10^-2= 1/100
.001= 10^-3= 1/1000
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32
Q

Kilo (k)

A

Thousand

10^3= 1,000

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33
Q

Milli (m)

A

One- thousandth

10^-3= 1/1000

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34
Q

Mega (M)

A

Million

10^6= 1,000,000

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35
Q

Micro

A

One-millionth

10^-6= 1/1,000,000

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36
Q

Giga

A

Billion

10^9

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37
Q

Tera

A

10^12

Trillion

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38
Q

Deci

A

10^-1= 1/10

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39
Q

Centi

A

10^-2= 1/100

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40
Q

Nano

A

10^-9

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41
Q

Pico

A

10^-12

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42
Q

Femto

A

10^-15

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43
Q

Atto

A

10^-18

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44
Q

F degree is _________ than C degree

A

Small

Difference is factor of 100/180= 5/9

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45
Q

32 degree F= ___________ C

A

0 degree Celsius

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46
Q

Conversion from F to C

A

C= (F-32) X 5/9

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47
Q

Conversion from C to F

A

F= C X 9/5 + 32

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48
Q

Conversion from C to Kelvin

A

K= C + 273

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49
Q

0 degree Celsius= _______ K

A
  1. 15 K
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50
Q

180 C=__________ K

A

372.15 K

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51
Q

Pressure:
1 atm=________ mmHG
1 atm= _______ torr

A

1 atm= 760 mmHG, 760 torr

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52
Q

1 mmHG= _______ torr

A

1 torr

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53
Q

760 torr= _______ atm

A

1 atm

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54
Q

1 atm= ________ psi
1 atm= ________ ft water
1 atm= ________ inches Hg
1 atm= ________ kPa

A
  1. 7 PSI
  2. 9 ft water
  3. 92 inches HG
  4. 325 kPa
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55
Q

Weight: 1 kg= ____ pounds

A

2.2 pounds

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56
Q

Weight: 1 pound=______ grams

A

454 grams

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57
Q

Volume: 1 liter=_______ ml

A

1000

58
Q

1 ml= ____ cc

A

1 cc

59
Q

1000ml=________ kg

A

1 kg

60
Q

1 liter= ____ quarts

A

1.06 quarts

61
Q

1 quart= ________ ml

A

946 ml

62
Q

20 drops= ________ cc
1 TB= _______ ml
1 ounce= ________ ml

A

1 cc
15 ml
29.6 ml

63
Q 
Distance:
1 cm= \_\_\_\_\_ inches
1 inch= \_\_\_\_\_\_ cm
1 kilometer= \_\_\_\_\_\_ miles
1 mile= \_\_\_\_\_\_\_\_ kilometers
A

1 cm= .39 inches
1 inch= 2.54 cm
1 kilometer= .62 miles
1 mile= 1.61 kilometers

64
Q

Don’t round off a calculated Quantity When all of the Numbers in your equation are _________________

A
  • Precise
  • If you can measure out 2.19 grams with your measuring device then do that
  • Know when it’s appropriate to be precise (weighing out NaCl) and when it’s okay to not be (baking a cake)
65
Q

When solving for Molarity what units should your answer be in?

A
  • Moles/liter= M

- if asked for a concentrated in mM, then you must convert from M to mM

66
Q

Molality

A

-Means of expressing the concentration of a certain species in solution in terms of moles per KG of water
-1 Molal solution: weight out MW in grams
-add exactly 1kg (1 Liter) of water.
Total volume will be greater than 1L

67
Q

Normality

A

-means of expressing the concentration of an acid in terms of hydrogen ion equivalents

68
Q

HCl= _________ normality

A

1 molar= 1 normality

Only one Hydrogen ion

69
Q

H3PO4= __________ Normality

A

1 molar= 3 normality

3 Hydrogen ions

70
Q

Equivalents

A

-Means of expressing the concentration of a species in terms of the number of ions released

71
Q

Equivalents:
100 mM of NaCl:
1 liter= ___________ mEq Cl-

A

1 liter= 100 mEq of Cl-

Because there is only one Cl- ion

72
Q

Equivalents:
100 mM of FeCl3
1 liter= _________ mEq of Cl-

A

1 Liter= 300 mEq of Cl-

Because there is 3 Cl- ions

73
Q

Osmolarity

A

-Means of expressing the overall number of particles in solution
Without regard to their identity
-Represents the total concentration of all species

74
Q

Osmolarity Example:

NaCL: 1 mM= _________ mOsm/liter

A

1mM= 2 mosm/liter

Na+ & Cl= 2 separate ions

75
Q

Osmolarity Example:

1 mM of MgSO4= ________ mOsm/liter

A

1 mM of MgSO4= 2 mOsm/liter

Mg 2+ & SO4-2

76
Q

1:1

A
  • Basic unit on which others are based

- 1 gram/ml

77
Q

1%

A

1 gram/100 ml (w/v)

1 ml/100 ml (v/v)

78
Q

Mg %

A
  • mg/DL

- mg/100 ml

79
Q

Definition of a Solution

A
  • Homogenous mixture in which one substance is dispersed in another as individual atoms or molecules
  • Liquid Solution:
    1. Transparent
    2. Passes through filter
    3. Does not settle upon standing
80
Q

Solvent

A

-Substance in which a solute dissolve to form a solution

81
Q

Solute

A

-Substance that dissolves in a solvent to form a solution

82
Q

Saturated Solution

A
  • solution in which solvent is unable to hold any more solute
  • solvent may be in equilibrium with an excess of undissolved solute
83
Q

Suspension

A
  • Finely divided solid particles dispersed in a liquid
  • Particles will settle upon standing under the influence of gravity
  • Must shake or mix before using
  • Particle size > 1 um trapped by filter
84
Q

Tyndall Effect

A
  • to distinguish solution from suspension which contains excess undissolved solute
  • particles not in solution scatter light
  • solvent will not be transparent to beam of light
85
Q

La Chatelier’s Principle

A

-if a system at equilibrium is perturbed, the system moves to counteract the effects of the change

86
Q

Crystalloid

A
  • Soltuion: particles are distributed as individual molecules
  • Particles are small enough to diffuse through a semi-permeable membrane
  • Particles are not trapped by a filter
87
Q

Examples of Crystalloids

A
  • LR
  • NS
  • D5W
88
Q

Colloid

A
  • Finely divided particles dispersed in a liquid (or gas)

- Particles are too small to settle under the influence of gravity

89
Q

Colloid

A
  • Particles are too large to diffuse through semi-permeable membrane
  • Particles not trapped by a filter
  • Particle size is 2-1,000 nm
90
Q

Examples of colloids

A
  • milk
  • dextran
  • hespan
  • albumin
  • fog
  • smoke
91
Q

Emulsion

A
  • Particles or globules of one liquid dispersed in another liquid
  • Liquids are immiscible
  • Layers may separate on standing
  • May contain emulsifier to prevent separation
92
Q

Micelle

A
  • Hydrophilic head
  • Aqueous solution
  • Hydrophobic tail
93
Q

Examples of emulsions:

A
  • Diprivian
  • TPN
  • Mayonnaise
  • Butter
  • Detergent on dirty clothes
94
Q

Soap

A

Ionic end= polar and hydrophilic
Hydro carbon Chain= non polar and hydrophobic
Cleaning action*

95
Q

Examples of suspensions:

A
  • NPH insulin
  • Maalox
  • Blood
  • Oral Antibiotics
96
Q

Density

A
  • mass per volume

- gram/ml

97
Q

Density of Water

A
  • maximum at 3.98 degree C
  • by definition, 1 gram/cc
  • 1 gram/ml
  • 1 kg/liter
98
Q

Specific Gravity

A
  • Weight of a substance compared to the weight of the same volume of a reference compound
  • Reference compound is usually water at its maximum density
  • Equivalent to density of substance in gram/ml without units
  • Ratio that has no units
99
Q

Ethanol:
Density= 0.789 grams/ml
Spec gravity=_____________?

A

Spec gravity= .789 (NO UNIT)

100
Q

Accuracy

A

Does the measurement give you the “right” answer

Think of darts and Bulls eye

101
Q

Reliability

A
  • Reproductibility

- extent to which a measurement gives the same answer every time

102
Q

Can a measurement be reliable but not accurate?

A

Yes
Example: tape measure if it’s been stretched
Example: non calibrated thermometer

103
Q

Precision

A
  • Extent to which a measurement can distinguish between 2 closely spaced values
  • Directly related to a number of significant digits in a measurement
  • Number of significant digits in value should be appropriate for the context: Measurment device, use of measurement
104
Q

Validity

A

-Extent to which a measurement provides a true indication of the parameter of interest
-A measurement can be accurate, reliable and precise but NOT valid
Example: IQ test as a measure of intelligence

105
Q

Questionable Validity

A
  • blood pH as a measure of acid/base status
  • pulse oximetry to monitor organ perfusion
  • pupillary size as an indicator of ICP
  • arterial line to monitor fluid status
106
Q

Gas

A
  • No definite shape or volume, conforms to container, fills container, flows and is easily compressible
  • gases that do NOT react chemically= homogenous mixture
  • Molcules high kinetic energy, large distances, do not interact with each other or container
  • Ideal gas= no intermolecular forces
107
Q

Liquids

A
  • Definite Volume
  • Conforms to shape of container
  • Flows, not compressible
  • Molecular distance small, medium kinetic energy, inter-molecular forces hold condensed state but allow molecules to slide against each other
108
Q

Solids

A
  • Definite volume, definite shape

- Molecules are densely packed, minimal kinetic energy, ordered structure

109
Q

Common Gases

A
  • O2
  • N2
  • CO2
  • CO
110
Q

Force

A

F= pressure/area

111
Q

Volume cylinder

A

V=area base X height

112
Q

Area of a circle

A

PiR(2)

113
Q

Volume of a cube

A

Side(3)

114
Q

Volume of a circle

A

4/3 Pi r (3)

115
Q

Temperature Conversions:
Boiling H20= F____=C______=K_________
Freezing H20= F______=C______=K________
O degree K=_________C=________F

A

Boiling H20= 212F= 100C= 373.15K
Freezing H20= 32F=0C-273.15K
O K= -273 C= -524 F

116
Q

1 atm=_____mmHg=_______ PSI=_____ kPa

A

1 atm= 760 mmHG= 14.7PSI=101 kPa

117
Q

Measuring Pressure

A
  • Atmospheric pressure exists
  • Gauge pressure is pressure within vessel (tank) ABOVE atmospheric pressure
  • Bourdon spring gauge (Pressure causes coiled copper to straighten) 02 tank gauges
118
Q

Avogadro’s Number and Hypothesis

A 
#= 6.02 X 10(23) the number of molecules in one mole of a substance
Hypothesis= one mole of gas at standard temperature (O degree C or 273.15K) and pressure (1 atm/1bar or 100 kPa) (STP) occupies a volume of 22.4 liters
-at equal temperature and pressure, equal volumes of gas contain equal numbers of gas particles
119
Q

How much volume is occupied by 2 moles of oxygen?

A

2 moles X 22.4L/mole= 44.8 liters of oxygen

120
Q

How much volume is occupied by a 2 mole of 50:50 mixture of O2/N2O?

A

2 moles X 22.4L= 44.8 Liters of O2/N2O

121
Q

Charles’ Law

A

-Volume of a gas is proportional to its absolute temperature as long as the pressure and amount of gas are held constant
VI/T1=V2/T2

NOTE: no p in this equation because pressure is constant
Temperature must be in Kelvin
Kelvin= C + 273.15

122
Q

Charles’ Law

A

-If you have a certain molar amount of gas and increase its temperature, the volume of space occupied by that gas will increase and vice versa.

123
Q

Boyles’ Law

A

-Volume of a gas is proportional to its pressure as long as the TEMPERATURE is held constant
P1V1=P2V2
Note: No T because temperature is constant
-decreased volume= increased pressure
-increased volume=decreased pressure
-Units of pressure and volume do not matter as long as they match on both sides

124
Q

Boyles’ Law:

What happens when we contract our diaphragm?

A
  1. Contract our diaphragm, the pressure inside the lungs drops below atmospheric pressure and air flows into our lungs
  2. When we release our diaphragm, pressure inside the lungs increase and volume of gas flows out of the lungs
125
Q

Gay-Lussac’s Law

A

-Pressure is proportional to absolute temperature (Kelvin) if volume is constant
-As temperature increases, pressure increases and vice versa
P1/T1=P2/T2
NOTE: no V because volume is constant

126
Q

The “Universal” or “Ideal” Gas Law

A

PV=nRT
Pressure (atm) X volume (L)= n (number of moles per gas) X R (universal gas constant (8.314 J/mol K)X T (temperature in K)

127
Q

Kinetic Molecular Theory: first 3?

A
  1. Ideal gases consist of a large number of tiny particles
  2. Gas particles are in constant, random motion
  3. Gas particles are small compared to the distance b/t them
128
Q

Kinetic Molecular Theory: 4-6?

A
  1. Volume of particle sis negligible
  2. Gas particles do not react chemically with the other container or each other
  3. No forces of attraction or repulsion b/t gas molecules or the container
129
Q

Kinetic Molecular Theory 7-9?

A
  1. Collisions b/t particles and container are perfectly elastic (no energy is lost)
  2. Pressure is caused by collisions b/t particles and container walls
  3. Average kinetic energy of a collection of gas particles depends on their absolute temperature and NOTHING else
130
Q

At STP (O degree C) and 1 atm (760mmhg) 1 mole of an ideal case will occupy how much space?

A
  1. 22.4 liters of space

At room temp (68 degree, 1 mole of ideal gas occupies 24L of space)

131
Q

Dalton’s Law of Partial Pressures

A

-The total pressure in a mixture of gases is equal to the sum of the partial pressures of the individual gases in the mixture
P (total)= PO2 + PCO2 + PN2O

The partial pressure of a gas is equal to the (Fi% X atmospheric pressure)
Need to know Fi% of the gas to make this law work

132
Q

What is the partial pressure of O2 in room air?

A
  • RA FiO2 is 21%
  • Atmospheric pressure is 760
    .21 X 760= 159.6 mmHG
    NOTE: pO2 and PCO2 is the partial pressure of Oxygen and CO2 in the blood.
133
Q

Composition of Dry Room Air

A

Nitrogen: 78%= 593 mmHG
Oxygen: 21%= 159 mmHG
Argon: 1%= 8 mmHG
CO2: .03%= .2 mmHG

134
Q

Anesthetic partial pressures: Iso, Sevo, Des

A
  1. 1% Isoflurance= 1 volume ISO/ 100 volume of carrier gas, Partial Pressure= 240 mmhg
  2. 2% Sevoflurance= 2 volumes SEVO/ 100 volume of carrier gas, Partial Pressure= 170 mmhg
  3. 6% Des= 6 cc des fluorane/100 cc oxygen/air, 669 mmhg
135
Q

Mole Fraction

A
  • Mole fraction of each component is the number of moles compared to the total number of moles of gas present
  • P1+P2= (N1+N2) RT/v

The partial pressure of each gas is directly related to the # of moles of the gas

136
Q

Law of LaPlace-Cylinders

A

T= P X R (T= wall tension, P= pressure, R= radius of cylinder)

  • As cylinder expands, the radius increases. As radius increases, wall tension increases.
  • As blood vessel dilates from increased fluid, wall tension in the vessel increases. Greater likelihood of rupture
  • Capillaries withstand high pressure better than large vessels especially veins because their radius is smaller and therefore the wall tension is lower
137
Q

Law of LaPlace- Spheres (alveoli)

A

T=Pr/2, or 2T= Pr, or P=2T/r
Alevoli without surfactant: P= 2T/r
-when R decreases, pressure in the alveoli increases if wall tension is the same
-When pressure increases in small alveoli (small R), P in small alveoli is higher than larger alveoli and air moves from small to large, and small alveoli collapse
-Result= atelectasis & ARDS. Premature infants born without surfactant

138
Q

Law of LaPlace- Spheres- Alveoli with surfactant

A

P=2T/r

  • surfactant allows for the surface tension to increase as radius increases
  • maintaining T keeps pressure within the alveoli the same
  • when P1=P2, gas does not move from A1 to A2
139
Q

Henry’s Law

A
  • amount of gas that dissolves in a liquid is directly proportional to the partial pressure of the gas in the gas phase over/around that liquid.
  • Blood is a liquid. O2, Co2, and Anesthetic agents=gas
140
Q

Oxygen Carrying Capacity Equation

A

CaO2= (1.34 X Hgb X SaO2) + 0.003 (PaO2)
-Answer is in ml/DL?

Note:
PAO2= alveolar oxygen partial pressure= 6X FiO2
PaO2= arterial oxygen partial pressure= 5X FiO2

141
Q

Graham’s Law

A

-Rate of diffusion through an office of a gas is inversely proportional to the square root of the molecular weight

R=1/ (square root of MW)
-Smaller molecules diffuse through membranes faster, larger slower
-Large molecules have higher molecular weight.
-This relates to second gas effect w/ Nitrous Oxide: rapid uptake of smaller molecule leaves higher concentration of large molecules that are left behind).
-this relates to diffusion hypoxia which can occur w/ N2O: reverse of
Second gas effect. Smaller molecules rush out and reduce concentration of the larger molecules.

Chem/Phys Exam 2 (1 decks)

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