This first chapter introduces a variety of basic concepts that are necessary to understand the advanced and complex nature of biochemistry. It provides a way to view matter and how it is composed (1.1), how to deal with measurements, units, uncertainty, significant figures and the scientific notation (1.2, 1.3 and 1.4) and finally explains some fundamental properties of matter (1.5).
1.1 - The Composition Of Matter
Matter is a physical substance that occupies space and has mass. It is often classified into two broad categories; pure substances and mixtures of pure substances.
1.1.1 - Pure Substances & Mixtures
Pure Substances
A pure substance consists of only one type of particle or chemical composition throughout. Its composition is uniform (i.e. the same everywhere) and doesn't change over time. There are two types of pure substances; elements and compounds.
- Elements are the most simple forms of matter. They cannot be broken down into simpler substances by chemical means. Each element is made of only one type of atom. Examples are oxygen (), gold () and carbon ().
- Compounds are substances formed when two or more elements are chemically combined in fixed proportions. The properties of a compound are different than those of its constituent elements. An example of a compound is water (), which consists of two hydrogen atoms and an oxygen atom in a 2:1 ratio. Other common examples are carbon dioxide () or sodium chloride ().
Mixtures
A mixture is a combination of two or more pure substances that are not chemically bonded together. Unlike compounds, the constituent substances of a mixture can be separated by physical means (i.e. separation techniques) such as filtration or distillation. A mixture can be classified based on how well the substances are mixed; homogeneous and heterogeneous mixtures.
- In homogeneous mixtures or solutions, the composition is uniform throughout. The different constituents are not visible, since the particles are mixed at a molecular level. Examples include saltwater (salt dissolved in water) and air (mixture of different gases such as nitrogen, oxygen, carbon dioxide, etc.).
- Heterogeneous mixtures are mixtures where the constituents are not uniformly distributed. It is often possible to visibly see the different substances. The composition may vary from one mixture to another. Examples include sand in water or granite.
Comparison
There are some key differences between pure substances and mixtures. They can be summarized as follows:
- Pure substances have a fixed composition, while mixtures can have variable compositions.
- Pure substances can be elements or compounds. Mixtures can be homogeneous or heterogeneous.
- Pure substances cannot be separated into other substances by physical methods, whereas mixtures can be separated into their individual components through separation processes.
1.1.2 - Separation Methods
To understand chemical processes, it is evident to study the properties of pure substances. As mentioned, before, these can be obtained from mixtures through a separation method. Common separation methods include:
- Filtration is used to separate an insoluble solid from a liquid. The mixture is passed through a filter which allows the liquid to pass through but traps the solid particles. An example is separating sand from water.
- Distillation can be used to separate liquids based on differences in boiling points. The mixture is heated and the constituent components with a lower boiling point will vaporize first. The vapor is then cooled and condensed back into a liquid. An example is separating alcohol from water.
- Evaporation separates a dissolved solid from a liquid. The liquid is evaporated, which leaves the solid behind. An example is recovering salt from saltwater.
- Chromatography is used to separate components based on differences in their movement through a medium. The mixture is applied to a medium (i.e. paper or gel) and a solvent is passed through. Different components move at different rates through the medium. An example includes separating pigments in ink.
- Centrifugation separates substances based on density differences. The mixture is spun at high speeds in a centrifuge, causing the denser components to move outward and separate from the lighter components. An example is separating blood into plasma and cells.
- Decantation can be used to separate a liquid from a heavier solid or immiscible liquid. It works by allowing the mixture to settle first. The liquid is then carefully poured off, leaving behind the solid or heavier liquid. An example is pouring oil off water.
- Magnetic separation separates magnetic materials from a mixture. A magnet is used to attract magnetic substances away from non-magnetic ones. An example is separating iron fillings from sand.
- Sublimation is used to separate a solid that sublimates (i.e. directly turns into gas) from other solids. Heat is applied to the mixture, causing the sublimable solid to vaporize, leaving behind the non-sublimable substance. An example is separating iodine crystals from salt.
- Crystallization can be used to separate a dissolved solid from a solution by forming crystals. The solution is evaporated or cooled, leading to the formation of solid crystals from the dissolved substance. An example includes obtaining pure sugar from a sugar solution.
1.1.3 - Properties
All pure substances have unique and consistent physical and chemical properties. These properties are fundamental to identifying and understanding substances in scientific study and application.
Physical Properties
Physical properties are characteristics that can be observed or measured without changing the substance’s chemical identity. Each physical property is consistent for a pure substance under specific conditions, allowing scientists to use these properties as benchmarks for identification. Examples of physical properties include:
- Density - The mass per unit volume.
- Melting point - The temperature at which a solid turns into a liquid.
- Boiling point - The temperature at which a liquid turns into a gas.
- Color - The appearance of a substance as perceived by the human eye (more specifically, the color response obtained by photoreceptor cells).
- Odor - The smell or scent of the substance.
- Solubility - The ability of a substance to dissolve in a solvent.
- Other include properties such as the refractive index, the thermal/electrical conductivity, viscosity, etc.
Chemical Properties
Chemical properties describe a substance's potential to undergo chemical changes that alter its composition. Chemical properties are particularly important in determining how substances interact and transform in various environments and reactions. In such a chemical reactions, substances can lose their chemical identities and form new substances with new physical and chemical properties. Some chemical properties include:
- Reactivity - How a substance chemically reacts with other substances.
- Flammability - The ability of a substance to ignite and burn in the presence of oxygen.
- Toxicity - The degree of which a substance can harm a living organism.
- Oxidation State - The potential of an element to lose, gain or share electrons in chemical reactions.
- Acidity/Basicity (pH) - The tendency of a substance to act as an acid or base, measured by pH.
- Others include properties such as radioactivity, redox potential, combustibility, etc.
It is important to note that pure substances can either be elements or compounds. While elements cannot be decomposed any further, compounds can. Chemical means allow to decompose these compounds back into their original elements.
1.2 - Measurements & The Metric System
It is more informative to describe matter in a quantitative way rather than a qualitative way. Quantitative properties (e.g. melting point, density, temperature, etc.) can be used to compile a profile that is unique. To create such a profile, measurements of these properties have to be carefully taken.
Def 1.1 - Measure
The size, capacity, extent, volume or quantity of anything, especially as determined by comparison with some standard or unit, is called a measure.
The definition of a measure suggests a need for a standard system of units. The system used in science and technology is the metric system. The version of this system that is being used today is called the “Système International d'Unités”, commonly known as the SI system. As the definition suggests, this system defines standard units for measurements.
Fundamental Quantity | Unit Name | Symbol |
length | meter | m |
mass | kilogram | kg |
temperature | kelvin | K |
time | second | s |
amount of substance | mole | mol |
electric current | ampère | A |
luminous intensity | candela | cd |
All other units that are used are so called derived units. Some examples include area (m), volume (m), density (kg/m) and velocity (m/s). Besides SI units, some other units will be used throughout this syllabus, since they are still in use in the scientific world today. These units can be found in the following table. They can of course, still be expressed in their equivalent SI units.
Quantity | Unit Name | Symbol | SI Definition | SI Equivalent |
length | Ångström | Å | 10m | 0.1 nanometer |
volume | liter | L | 10m | 1 decimeter |
energy | calorie | cal | (kg m s) | 4.184 joules |
One of the advantages of the metric system is that all basic units are multiplied or divided by multiples of ten. This allows mathematical manipulation to be simple, by moving the decimal point left or right. The next table lists how these multiples of ten are prefixed and what symbol is used to denote them.
Multiple or Part of Ten | Prefix | Symbol |
1,000,000,000 | giga- | G |
1,000,000 | mega- | M |
1,000 | kilo- | k |
100 | hecto- | h |
10 | deka- | da |
0.1 | deci- | d |
0.01 | centi- | c |
0.001 | milli- | m |
0.000001 | micro- | |
0.000000001 | nano- | n |
0.000000000001 | pico- | p |
1.3 - Uncertainty & Significant Figures
It is unlikely that a series of measurements of the same property of the same object will result in precisely the same value. It is unavoidable to make a “best guess” estimate. This estimate is called the uncertainty or variability. The difference between the true (unobtainable) value and the measured value is the error. To deal with these variabilities, scientists use a system that takes advantage of a so called significant figures. The system eliminates the need for a notation and indicates the uncertainty by means of the number of digits that are used instead:
In the small example above, the first notation has 1 significant figure, the second has 2, and so forth. In this way, degrees of uncertainty are communicated through the numbers of significant figures. Here, it is important to distinguish when the number 0 is significant or not. The following bullet points summarize the importance of the number 0 when working with significant figures:
- A trailing zero is significant (e.g. 4.130)
- A zero within a number is significant (e.g. 35.06)
- A zero before a digit is not significant (e.g. 0.082)
- A number ending in zero with no decimal point is ambiguous (e.g. 20)
1.4 - Scientific Notation
The scientific notation often called the exponential notation is a convenient method for preventing ambiguity in the reporting of measurements and for simplifying the manipulation of very large and very small numbers. To express a number in scientific notation, one writes the number between 1 and 10 multiplied by 10, raised to a whole-number power.
Example 1.1 - Scientific Notation
Here are some examples on how to write numbers in scientific notation.
- 3 is written:
- 24 is written:
- 346 is written:
- 2537 is written:
In general, the following steps can be used to transform a number to its scientific notation:
- For any number greater than 1, the decimal is moved to the left to create a coefficient between 1 and 10.
- Next, an exponential factor is created, whose positive power is equal to the number of places that the decimal point had to be moved to the left.
The process is somewhat different for numbers that are smaller than 1. The following algebraic rule can be used:
- For any number less than 1, the decimal is moved to the right to create a coefficient between 1 and 10.
- Next, an exponential factor is created, whose negative power is equal to the number of places that the decimal point had to be moved to the right.
1.5 - Fundamental Properties Of Matter
Matter has a variety of properties that can be measured. One of the most useful of these properties is mass, the reason being that it stays constant with environmental changes such as fluctuations in temperature or pressure.
Mass & Volume
Def 1.2 - Mass
Mass is a measure of the quantity of matter.
It is important to note that mass and weight are not the same thing! A mass has a weight, because it is under the influence of a gravitational field. For example, you weigh less on the moon, but your mass does not change.
A sample of a pure substance or mixture has a mass. Besides mass, this sample also takes up space. This space is called the volume.
Def 1.3 - Volume
Volume is the amount of space that a sample occupies.
Density
Other useful properties are necessary. For example, the characterization of substances requires measurements of their physical properties (e.g. mass). This has to be done without changing the substance's nature. One of these properties is density.
Def 1.4 - Density
Density is a property of substances in any physical state (i.e. gaseous, liquid or solid). It is defined as the mass per unit volume.
Density is a property that can be used for a variety of actions, such as:
- Evaluating the purity of solids and liquids.
- Estimating the amount of dissolved solids in solutions.
- It acts as a conversion factor for translating mass into volume and vice versa.
Because volume changes with temperature, density decreases. Therefor, a measurement of density should always be accompanied by a measurement of temperature.
Temperature
Some substances are cold, others are hot. We say that substances can gain or lose heat, depending on whether they are cooler or hotter than the environment. Heat is often describes as temperature.
Def 1.5 - Temperature
Temperature is a measure of the heat of a substance.
Temperature is measured in Kelvin (K), which is an SI unit. The freezing point of water, 0°C is about 273.15K. For convenience this value is often rounded to 273, hence:
Finally, heat always "flows" from hotter regions to colder regions. In other words, heat always moves in the negative direction on a temperature gradient. This can be compared to water, which will always flow downhill.
Heat & Calorimetry
Heat is a form of energy. Focusing on loss or gain of heat, it is possible to enlarge the list of properties useful for the characterization of different substances. Each substance has a different capacity to absorb heat. This capacity can be measured by noting the rise on temperature of a fixed mass of a substance that has absorbed a known amount of heat. The question now remains on how to measure heat (i.e. we know how to measure temperature and mass). The non-SI unit of heat is the calorie (cal). The official SI unit of heat is called the joule (J).
Example 1.2 - Heating of
The temperature of 1.0g of water rises 1°C when 1 cal of heat is absorbed. In other words, when 1.0g of water absorbs 4.184J of energy, the temperature will rise 1°C.
The characteristic response (increase in temperature) of a given mass of a given substance exposed to a given amount of heat is expressed by a quantity called the specific heat (). If the mass is in grams and the temperature on the Celsius scale, and the heat in joules, the units of specific heat are:
Specific heat should not be confused with the heat capacity. Heat capacity is the capacity of a given sample to absorb heat and therefore depends not only on the substance's characteristics, but also on its mass. The greater the mass, the greater the amount of heat that the sample can absorb.
Continue reading →
Chapter 2 - Atomic Structure
Sources
- “General, Organic, and Biochemistry” (2nd Edition) - Blei I. and Odian G.
No sources other than the syllabus where used to write this chapter.