Introduction to Organic and Biochemistry (CHE 124)

1
Introduction to Organic and Biochemistry(CHE 124)

• Reading Assignment
• General, Organic, and Biological Chemistry An
  Integrated Approach
• 4rd. Ed. Ramond
• Chapter 1
• Science and Measurements
• Answers to odd numbered problems in textbook are
  found in the books index.

2
What is Chemistry?

• Chemistry - the study of matter and the changes
  that it undergoes (e.g. reactions).
• Chemistry is the central science. It unifies the
  sciences in biology, physics, engineering,
  medicine, pharmacy, etc.

3
What is the scientific Method?

• The scientific method is a way of gathering and
  interpreting information about chemistry (see
  next slide).
• Hypothesis
• tentative explanation (educated guess) for
  observations and known facts.
• Theory
• an experimentally tested explanation of an
  observed behavior. (a well tested hypothesis)
• Law
• statement that describe things that are
  consistently and reproducibly observed. (a well
  tested theory)

4
Scientific Method
5
What is Matter?

• Matter
• anything that has mass and occupies space.
• Weight measure of gravitational pull against
  matter.
• Mass measure of amount of material.
• Phases of Matter
• Solids
• Fixed volume and shape
• Liquids
• Fixed volume, indefinite shape.
• A liquids takes on the shape of the container.
• Gases
• Indefinite shape and volume.
• A gas takes on both the shape and volume of the
  container.

6
Properties of Substances

• Every pure substance has its own unique set of
  properties that serve to distinguish it from all
  other substances.
• Look them up in the chemical literature
• These properties are classified as
• Intensive properties independent of amount of
  substance. ( E.g. m.p., b.p., density)
• Extensive properties dependent on amount of
  substance. (e.g. mass, volume)
• Chemical properties observed when the substance
  takes part in a chemical reaction
• becomes a new substance
• E.g. Heating Mercury (II) oxide to produce
  oxygen.
• Does the substance react with oxygen?
• Physical properties
• No chemical change is required
• Examples
• Melting point (m.p.) temp. when substance
  changes from solid to liquid.
• Boiling point (b.p.) temp. when substance
  changes from liquid to gas.
• Density
• Solubility amount of solute that dissolves in a
  give amount (100g) of solvent at a specific temp.
• Color

7
What is Energy?

• Energy
• The ability to do work and / or to transfer heat.
  
• Potential energy
• stored energy
• Kinetic energy
• energy of motion

8
Measurements

• Chemistry is a quantitative science. We use the
  SI system of measures.
• SI Units International System of Measure
• Common name the metric system
• Based on the decimal (powers of ten)
• Kg, L, K, C
• English system is used in the United States.

9
Measuring Length

• SI unit of Length meter (m)
• Definition of meter - the distance light travels
  in 1/299,792,458 of one second
• 1 m 39.37 in.
• English units
• mile (m), yard (yd.), foot (ft.),inch (in.)
• Instruments used to measure length
• Meter stick
• Micrometer

10
Measuring Volume

• SI Unit of Volume Liter (L)
• Volume is derived from SI unit of length.
• Units of volume
• cubic meter (m3) 1000 L
• cubic centimeter (cm3 or cc) milliliters (mL)
• English Units
• Gallon (gal.), quart (qt.), pint (pt.), cup (c),
  teaspoon (tsp.) , table spoon (tbsp.), fluid
  ounce (oz.)
• Instruments used to measure volume
• Graduated cylinder, pipet or buret, digital
  micropipet

11
Measuring Mass

• SI Unit kilogram (kg)
• Definition of kg
• The kilogram is the unit of mass it is equal to
  the mass of the international prototype of the
  kilogram.
• Kilogram (kg) 1000 g
• 1 gram (g) 1000 mg
• English Units
• Ton (ton.), pound (lb.), ounce (oz.)
• Instruments used to measure mass
• Balance
• Scale

12
Mass vs Weight

• Mass
• Amount of matter in a sample.
• Weight
• The effect of gravity on the matter.

13
SI Units
Base Quantity Name Symbol
Length Meter m
Mass Kilogram kg
Time Second s
Electric current Ampere A
Temperature       Kelvin K
Amount of substance Mole mol
Luminous intensity Candela cd
Source http//physics.nist.gov/cuu/Units/units.ht
ml
14
Derived SI Units
Derived Quantity Name Symbol
Area square meter m2
Volume cubic meter m3
Speed, velocity meter per second m/s
Acceleration meter per second squared   m/s2
15
Derived SI Units with Special Names
Derived quantity Name Symbol   Expression  in terms of  other SI units Expressionin terms ofSI base units
force newton N   - mkgs-2
pressure, stress pascal Pa N/m2 m-1kgs-2
energy, work, quantity of heat joule J Nm m2kgs-2
power, radiant flux watt W J/s m2kgs-3
Celsius temperature degree Celsius C   - K
16
Metric Prefixes
Prefix Symbol Multiple Multiple
mega M 1,000,000 1 x 106
kilo k 1,000 1 x 103
hecto h 100 1 x 102
deca da 10 1 x 10
Unit ------- 1 1
deci d 0.1 1 x 10-1
centi c 0.01 1 x 10-2
milli m 0.001 1 x 10-3
micro µ 0.000001 1 x 10-6
nano n 0.000000001 1 x 10-9
17
English Conversions
Length Volume
1 mile (m) 5280 feet (ft.) 1 gallon (gal) 4 quarts (qt.)
1 ft. 12 inches (in.) 1 qt. 2 pints (pt.)
1 yard (yd.) 3 ft. 1 pt. 2 cups (c.)
1 pt. 16 fluid ounces (fl. oz.)
Mass 1 c. 8 fl. oz.
1 ton 2000 pounds (lbs.) 1 in3 16.387 cm3
1 lb. 16 ounces (oz.) 2 tablespoons (T or tbsp) 1 fl. oz

18
English to Metric Conversions
Length
1 in. 2.54 cm (exact)
1 m 39.37 in
1 mi 1.609 km
Volume
1 ft3 28.32 L
1 L 1.057 qt
1 gal. 3.785 L
1 tsp. 5 mL
1 tbsp 15 mL 1 mL 15 drops (gtt)
Mass
1 lb. 453.6 g 0.4536 kg
1 g 0.03527 oz.
1 kg 2.2 lbs.
19
Typical Conversions Problems

• The dosage on a bottle of medicine reads Take 2
  tablespoons every twelve hours. Convert this
  volume to mL.
• The box at the pet store states Aquarium volume
  is 55 gal.. Convert this volume to liters.
• The distance (length) from Clinton to Vicksburg
  is appr. 32 miles. What is this distance in cm?
• A marathon is defined as 42.195 km. What is this
  distance in miles?

20
Measuring Temperature

• Factor that determines the direction of heat flow
• SI units C or K
• Fahrenheit (F)
• Named after German instrument maker Daniel
  Fahrenheit (1686 -1736)
• Celsius (C) (old name centigrade)
• Named after Swedish astronomer Anders Celsius
  (1701-1744)
• Kelvin (K)
• Do not use () or degree in relation to K.
• Defined as 1/273.16 of the difference between the
  lowest attainable temp. (0K) and the triple point
  of water (0.01 C)
• Instruments to measure
• Mercury thermometer
• Mercury expands and contracts as temperature
  changes. Tube contains only 2 of the Hg in
  thermometer.
• Digital thermometer
• Water boils 212 F 100 C 373.15 K
• Water freezes 32 F 0 C 273.15 K
• Triple point is the temp. /pressure combination
  at which water is capable of coexisting as a
  solid, liquid and gas.

21
Converting Temperature Scales

• Comparing F to C
• 0 C is 32 F
• Freezing point of water
• 100 C is 212 F
• Boiling point of water
• There are 180 Fahrenheit degrees for every 100
  Celsius degrees, so each C is 1.8 times larger
  than each F
• Comparing C to K
• Celsius degree and Kelvin degree are the same
  size.

22
Scientific Notation

• A way of dealing with very large or very small
  numbers.
• Show examples on the board
• Avogadro number 6.022 X 1023
• 83,000
• 0.000056
• See table 1.3 p.12

23
Accuracy vs Precision

• Accuracy
• How close a reported value is to the real value.
  (see next slide about error).
• Precision
• A measure of how close repeated measurements are
  to one another.

24
Uncertainties in Measurements

• How much solution is in the large graduated
  cylinder?
• How much solution is in the small graduated
  cylinder?
• Do these two measurements have the same
  uncertainty?

25
Uncertainties in Measurements

• Three volume measurements with their
  uncertainties
• Large graduated cylinder, 8 1 mL
• Small graduate cylinder, 8.0 0.1 mL
• Pipet or buret, 8.00 0.01 mL
• To denote how much uncertainty is in a
  measurement, Significant figures are used.
• Significant Figures
• Every measurement carries uncertainty
• All measurements must include estimates of
  uncertainty with them
• There is an uncertainty of at least one unit in
  the last digit
• Text convention
• Uncertainty of in the last digit is assumed but
  not stated

26
Significant Figures

• Significant figures are meaningful digits in
  measurements
• In 8.00 mL, there are three significant figures
• In 8.0 mL, there are two significant figures
• In 8 mL, there is one significant figure

27
Ambiguity in Significant Figures

• Consider the measurement, 500 g
• If the measurement was made to the nearest 1 g,
  all three digits are significant
• If the measurement was made to the nearest 10 g,
  only two digits are significant
• Resolve by using scientific notation
• 5.00 X 102 g
• 5.0 X 102 g
• See Table 1.5 p. 17

28
Rounding

• Rounding off numbers
• If the first digit to be discarded is 5 or
  greater, round up
• If the first digit to be discarded is 4 or
  smaller, round down

29
Significant Figures in Calculations

• Addition and Subtraction
• Count the number of decimal places in each number
• Round off so that the answer has the same number
  of decimal places as the measurement with the
  greatest uncertainty (i.e., the fewer number of
  decimal places).
• Multiplication and Division
• When multiplying or dividing two numbers, the
  answer is rounded to the number of significant
  figures in the less (or least in the case of
  three or more) measurements
• 2.40 X 2 5
• Exact Numbers
• Exact numbers carry an infinite number of
  significant figures
• Exact numbers do not change the number of
  significant figures in a calculation
• Example The numbers 1.8 and 32 in the conversion
  between Fahrenheit and Celsius

30
Dimensional Analysis / Factor Label / Converting
Units

• In many cases throughout your study of chemistry,
  the units (dimensions) will guide you to the
  solution of a problem
• A correct answer must have the NUMBER and UNITS!
• Conversion factors are used to convert one set of
  units to another
• Only the units change
• Conversion factors are numerically equal to 1
• 1L 1000 cm3 Choose a conversion factor that
  puts the initial units in the denominator
• The initial units will cancel
• The final units will appear in the numerator

31
Some Examples

• Convert 25 mL to L.
• Convert 200 pounds to grams.
• Convert 20 miles to kilometers.
• Convert 25 microliters to liters.

32
Properties of Substances

• Every pure substance has its own unique set of
  properties that serve to distinguish it from all
  other substances.
• Look them up in the chemical literature
• These properties must be intensive.
• Intensive properties independent of amount of
  substance. ( E.g. m.p., b.p., density)
• Extensive properties dependent on amount of
  substance. (e.g. mass, volume)
• Chemical properties observed when the substance
  takes part in a chemical reaction
• becomes a new substance
• E.g. Heating Mercury (II) oxide to produce
  oxygen.
• Does the substance react with oxygen?
• Physical properties
• No chemical change is required
• Examples
• Melting point (m.p.) temp. when substance
  changes from solid to liquid.
• Boiling point (b.p.) temp. when substance
  changes from liquid to gas.
• Density
• Solubility amount of solute that dissolves in a
  give amount (100g) of solvent at a specific temp.
• Color

33
Density

• The density of a substance is its mass divided by
  its volume. (the amount of mass contained in a
  given volume.)
• Units g / mL
• Density of
• Water is 1 g / mL
• Temperature must be stated.
• density changes with changes in temperature.
• Note table 1.8 p. 22

34
Specific Gravity

• Relates density of a substance to that of water.
• Measured using refractometeres, hydrometers or
  test strips.
• Used to determine
• acid level in car batteries
• antifreeze level in car radiators
• alcohol content in beer and wine
• Urine to diagnose kidney problems
• Temperature must be specified since the density
  of the substance and water vary, but not
  necessarily at the same rate.

Specific Gravity Density of substance 0.785
g/mL 0.785 Density of water
1.00 g/mL
35
Specific Heat

• Relates energy (in calories), mass (in grams),
  and temperature (in degrees Celsius).
• Units cal / g C
• Relates the mass, temperature, and energy.