Chapter 1 The Language of Chemistry

Topics to be covered:  matter and properties; measurements; significant figures; Calculations; units; and heat and energy

What is Chemistry?
Chemistry deals with the structure of matter and changes that matter (the “stuff” that things are made of)  undergoes.

All matter consists of chemicals
—things you find in food, cloths, house, cars, medicines, vitamins, sugar, table salt,...and toxic and hazardous substances, natural toxins, etc.

Why study chemistry?

• To understand the things, the phenomena, the processes we encounter in living systems and in the material world (e.g., protease inhibitors for HIV and cancer treatments, DNA cleavage for cancer treatment)

• To solve the problems you may encounter in life science, health fields, everyday-life, and material world
(e.g., drug development, disease control and treatment, pollutions)

**E.g.,  The cover story:  Sickled cells are formed because the glutamate at position 6 of the b chain of hemoglobin is mutated to a valine which causes coagulation of hemoglobin.

—consist of two or more pure substances, such as milk, blood, wood...

Homogenized milk contains 4% butterfat.  Low-fat milk contain 2% or 1% butterfat, which is 50% or 75% reduction of the butterfat relative to homogenized milk.

Homogeneous mixture—completely uniform, and can be mixture of materials with different physical states.

 gas/gas:  air
 liquid/gas:  moist air, soda water, (oxygenated water)
 liquid/liquid:  vinegar, wines
 Solid/liquid:  Brine (salt water)
 Solid/solid:  metal alloy (mercury amalgam filling)

Heterogeneous mixture—not uniform

 Soil; orange juice with pulps; milk; blood; etc.

Physical properties and physical changes

 Physical property:  quality and condition of a substance that can be observed without changing the substance's composition, such as color, solubility, odor, conductivity, magnetism, melting temperature...

Examples:  change of physical states (e.g., melting, evaporation, and sublimation)

Chemical reactions and chemical properties

 Chemical reactions:  changing of the composition of matter, such as the burning of gasoline, digestion of food...

  iron + sulfur —(heat)--> iron sulfide

 gasoline + oxygen ---> carbon dioxide + water
  (reactants)                           (Products)

 Chemical property:  A substance's behavior in chemical reactions
 (observation:  change in color or/and order?)
The Scientific Method
—a systematic and logic approach to the solution of problems (The method that leads to all scientific discoveries, such as the discovery of penicillin, enzymes are proteins, etc.).

An example is also given in the textbook on:  "Role of DNA in living systems"

1)  Observation—What happens?

2)  Hypothesis—Possible causes?  Reasonings?

3)  Experiments—Confirming or revising hypothesis

4)  Theories—Hypotheses supported by the results of  experiments
Scientific Laws:  Processes or phenomena, to which there is no exception

e.g., The law of conservation of mass and energy in chemical reactions

 —A part of daily life
  Examples:  cooking, car mileage, utility bills, ...

 —Fundamental of research
  Examples:  How fast reactions go; how large binding affinities are; molecular structure (inter-atomic distances and angles), etc.

 —In health sciences
  Examples:  blood pressure, hemoglobin content, blood sugar level, alcohol level, pharmokinetics, drug pathways, metabolite concentrations, etc.

Qualitative measurement
 —evaluation without expression as numbers
 Examples:  Blood pressure is "HIGH"; "very high", "above average", or...

Quantitative measurement
 —evaluation giving results as numbers
 Examples:  Blood pressure is "140/100"; and...

Measures must be accurate and precise to be useful.

Accuracy:  deviation from the true value; the smaller the deviation, the higher the accuracy
 Example:  measurement of water melting temperature  at 1 atm
 Accuracy (from high to low):
 0.01 °C > 0.5 °C > –1.0 °C > –1.5 °C > 2.0 °C
When there are several measurements, the accuracy of the measurements is determined by the average or the mean value.

Precision:  reproducibility of different measurements (can be determined with “standard deviation”, SD)
 Experiment 1 (°C):
0.020, 0.050, 0.030, –0.020, 0.000, 0.020, 0.030, –0.010

 Experiment 2 (°C):
0.020, 0.120, 0.30, 0.110, –0.100, 0.210, –0.150, –0.200

 Experiment 3 (°C):
1.020, 1.050, 1.030, 0.980, 0.990, 1.030, 1.010, 1.020, 0.990

Experiment 1 has very high accuracy and high precision (small SD), average = 0.015 and SD = 0.022; 2 has lower accuracy and lower precision (high SD), average = 0.039 and SD = 0.166; and 3 has the lowest accuracy but high precision (low SD), average = 1.013 and SD = 0.022.

In the above examples, you ought to be able to tell very easily without a calculator which one has high precision and which one has high accuracy.  However, the statistics mode on a calculator can provide you with quantitative comparison.

Systems of measurement

Metric system—based on multiples of 10.  (The most useful system in scientific research)

SI unit (International System of Units):
Base units:  meter (m); kilogram (kg); second (s, sec); Kelvin (K); mole (mol); etc.

Some commonly used non-SI units:  cm, L, gm, °C, atm, mm/Hg, cal,

Scientific notation:  used for very large and small numbers

Examples:  1 mole of atoms/molecules = 6.02 x 1023; a chemical bond ~ 2 x 10–8 cm; distance between the sun and the earth ~ 1.5 x 108 km; size of a bacterium ~ 0.5-2 x 10–6 m;

Significant figures:  all the digits known for certain plus one estimated digit

Example:  Mercury thermometer
37                38                 39                40 (°C)
Correct readings:  37, 38, 39, and 40 °C (to the first digit)
 Estimate digit:  0.X °C (to the 10ths)

37.0     37.5       38.0        38.5        39.0 (°C)
 Correct readings:  37.X °C, etc. (to the 10th digit)
 Estimate digit:  0.0X °C (to the 100ths)

1) Nonzero digit are significant
     154 m, 2.34 m, 1.96 x 104 m —all 3 significant figures
2) Zeros between nonzeros are significant.
     5706 m, 4.002, 1.102 x 10–2 —4 significant figures
3) leading zeros are not significant
     0.0034, 0.000067, 0.86 —2 significant figures
4) zeros after a decimal point at the end of a number are significant
     6.00 m, 0.0450 m and 15.0 x 103 m—all have 3 significant figures

What's the difference between 5 kg and 5.000 gm? 37 °C and 37.0 °C?  1 m and 1.00 m?
Significant Figures in Calculations

What's "wrong" with the following calculation of three measurements?
    1.2 g + 1.1 g + 1.23 g + 2.002 g = 5.532 g

Anything wrong with the following "addition"?
    1.2 + 1.1 + 1.23 + 2.002 = 5.532

What's "wrong" with the following calculation ?
    1.2  m x 1.1 m = 1.32 m2

Anything wrong with the following "multiplication"?
    1.2 x 1.1 = 1.32

Addition and subtraction:
The result can have no more digits to the right of the decimal point than the measurement with the least number of decimal places.

Multiplication and division:
The result cannot have more significant figures than the measurement with the smallest number of significant figure.

Correct way of using a scientific calculator— Don't take all the digits from the calculator, but need to consider the significant figures.

 Length and Volume
metric system:
kilometer — 1 km = 103 m
micrometer — 1 mm = 10–6 m (Unit for measuring micro-organisms)
nanometer — 1 nm = 10–9 m (Unit for measuring wavelength of visible lights, ~350-750 nm)

Å — 10–8 cm (A common unit for measuring micro-organisms, size of molecules, and intermolecular distance)

Volume—Space occupied by matter (liter, 1 L = 1000 cm3; cm3 = mL = cc)

1 gallon ~ 4 L

Converting units
Examples:  The recommended dosage of a certain drug is 5.0 mg of the drug for each kg of the body weight.
(a) How many mg of this drug should an 85-kg person receive?
(b) Express this dosage in g and mg.

 Mass—quantity of matter an object contains, measured by comparing it to a standard mass of 1 kg.
    1g = the mass of 1 cm3 water at 4 °C

(Weight is not equivalent to mass, since weight changes with conditions and is a measure of the force the gravity pulls.  What is your body mass and body weight?)

        (Check here if you want to learn more about mass and weight.)

Density—ratio of the mass of an object to its volume (amount of mass per unit volume of an object)
        (What is the similarity with the definition of concentration?)

Specific gravity—ratio of the density of a substance to the density of a reference substance at the same temperature (relative density of one substance to another)

What the difference between the two in terms of the measured values?


(Qualitative expression:  hot, warm,...; cold, freezing...)

Quantitative determinations:  temperature scales

1) Fahrenheit scale by Daniel Fahrenheit:
 freezing point of water = 32 °F
 boiling point of water = 212 °F
 (There are 212 – 32 = 180 °F between these two points.)

2) Celsius scale after Anders Celsius:
 freezing point of water = 0 °C
 boiling point of water = 100 °C
 (There are 100 °C between these two points.)

The relationship between these two scales is as below (or, you may think about in terms of proportionality).
  (°F – 32)/180  = °C/100   or  180/100  =   (°F – 32)/°C

The equations can be converted to
  °C = (5/9)(°F – 32)    or

  °F = (5/9)°C + 32

Average normal oral temperature:  98.6 °F (37 °C)
With a high fever:  106 °F (~41 °C)

(Different scales are used in different part of world.)

3) Kelvin scale (K, not °K) after Lord Kelvin:
  K = °C + 273        (This is the most important temperature scale used in physical sciences.)

 (freezing point of water = 273 K, boiling point of water = 373 K)
Heat and Heat Capacity

1 Calorie (cal)—quantity of heat to raise the temperature of 1-g pure water by 1 °C .

Although cal is commonly used in chemistry, the SI unit joule (J) is even more commonly used.  Try to do some conversions based on 1 cal = 4.184 J.

Specific heat capacity—quantity of heat required to raise the temperature of a substance by 1 °C, its unit is thus cal/(g x °C ).
This is a physical property of a substance.

heat (cal) = specific heat (cal/g °C ) x mass (g) x temperature change (°C)

The following relationships maybe helpful to get a better idea about the above equation:
If heat and mass are constants (e.g., fixed amounts), the larger the specific heat, the smaller the change in temperature.
If the specific heat and the temperature change are constants, the larger the mass, the greater amount the mass can absorb.
If specific heat and mass are fixed, the more heat provided and larger the change in temperature.

You should have better idea about the following questions:
What's the relative specific heat between cotton and aluminum, and between steel and the ceremic tiles on the space shutter? 

1) How much heat (in cal) is needed to raise the temperature of 11 g ice from –15 to –5 °C?  [Specific heat of ice = 0.5 cal/g•°C]
        Heat = ???
2) When you mix 5.0 g of water at 20 °C with 10 g of water at 50 °C, what is the final temperature?
      (After you think about the question, and work out your thought, you may want to check the answer and explanation here.)