Ross and Wilson Anatomy and Physiology in Health and Illness: With access to Ross & Wilson website for electronic ancillaries and eBook, 11 ed.

2. Introduction to the chemistry of life

Because living tissues are composed of chemical building blocks, the study of anatomy and physiology depends upon some understanding of biochemistry, the chemistry of life. This chapter introduces core concepts in chemistry that will underpin the remaining chapters in this book.

Atoms, molecules and compounds

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Learning outcomes

After studying this section, you should be able to:

image define the following terms: atomic number, atomic weight, isotope, molecular weight, ion, electrolyte, pH, acid and alkali

image describe the structure of an atom

image discuss the types of bond that hold molecules together

image outline the concept of molar concentration

image explain the importance of buffers in the regulation of pH.

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The atom is the smallest unit of an element that exists as a stable entity. An element is a substance containing only one type of atom, e.g. iron contains only iron atoms. When a substance contains two or more different types of atom, it is called a compound. For instance, water is a compound containing both hydrogen and oxygen atoms.

There are 92 naturally occurring elements, but the wide variety of compounds that make up living tissues are composed almost entirely of only four: carbon, hydrogen, oxygen and nitrogen. Small amounts (about 4% of body weight) of others are present, including sodium, potassium, calcium and phosphorus.

Atomic structure

Atoms are mainly empty space, with a tiny central nucleus containing protons and neutrons surrounded by clouds of tiny orbiting electrons (Fig. 2.1). Neutrons carry no electrical charge, but protons are positively charged, and electrons are negatively charged. Because atoms contain equal numbers of protons and electrons, they carry no net charge.

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Figure 2.1 The atom, showing the nucleus and four electron shells.

These subatomic particles differ also in terms of their mass. Electrons are so small that their mass is negligible, but the bigger neutrons and protons carry one atomic mass unit each. The physical characteristics of electrons, protons and neutrons are summarised in Table 2.1.

Table 2.1 Characteristics of subatomic particles

Particle

Mass

Electric charge

Proton

1 unit

1 positive

Neutron

1 unit

neutral

Electron

negligible

1 negative

Atomic number and atomic weight

What makes one element different from another is the number of protons in the nuclei of its atoms (Fig. 2.2). This is called the atomic number and each element has its own atomic number, unique to its atoms. For instance, hydrogen has only one proton per nucleus, oxygen has eight and sodium has eleven. The atomic numbers of hydrogen, oxygen and sodium are therefore 1, 8 and 11 respectively. The atomic weight of an element is the sum of the protons and neutrons in the atomic nucleus.

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Figure 2.2 The atomic structures of the elements hydrogen, oxygen and sodium.

The electrons are shown in Figure 2.1 as though they orbit in concentric rings round the nucleus. These shells diagrammatically represent the different energy levels of the electrons in relation to the nucleus, not their physical positions. The first energy level can hold only two electrons and is filled first. The second energy level can hold only eight electrons and is filled next. The third and subsequent energy levels hold increased numbers of electrons, each containing more than the preceding level.

The electron configuration describes the distribution of the electrons in each element, e.g. sodium is 2 8 1 (Fig. 2.2).

The chemistry of life depends upon the ability of atoms to react and combine with one another, to produce the wide range of molecules required for biological diversity. The atomic particles important for this are the electrons of the outermost shell. An atom is reactive when it does not have a stable number of electrons in its outer shell, and may donate, receive or share electrons with one or more other atoms to achieve stability. This will be described more fully in the section discussing molecules and compounds.

Isotopes

These are atoms of an element in which there is a different number of neutrons in the nucleus. This does not affect the electrical activity of these atoms because neutrons carry no electrical charge, but it does affect their atomic weight. For example, there are three forms of the hydrogen atom. The most common form has one proton in the nucleus and one orbiting electron. Another form (deuterium) has one proton and one neutron in the nucleus. A third form (tritium) has one proton and two neutrons in the nucleus and one orbiting electron. These three forms of hydrogen are called isotopes (Fig. 2.3).

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Figure 2.3 The isotopes of hydrogen.

Because the atomic weight of an element is actually an average atomic weight calculated using all its atoms, the true atomic weight of hydrogen is 1.008, although for most practical purposes it can be taken as 1.

Chlorine has an atomic weight of 35.5, because it contains two isotopes, one with an atomic weight of 35 (with 18 neutrons in the nucleus) and the other 37 (with 20 neutrons in the nucleus). Because the proportion of these two forms is not equal, the average atomic weight is 35.5.

Molecules and compounds

It was mentioned earlier that the atoms of each element have a specific number of electrons around the nucleus. When the number of electrons in the outer shell of an element is either the maximum number (Fig. 2.1), or a stable proportion of this fraction, the element is described as inert or chemically unreactive, i.e. it will not easily combine with other elements to form compounds. These elements are the inert gases – helium, neon, argon, krypton, xenon and radon.

Molecules consist of two or more atoms that are chemically combined. The atoms may be of the same element, e.g. a molecule of atmospheric oxygen (O2) contains two oxygen atoms. Most molecules, however, contain two or more different elements, e.g. a water molecule (H2O) consists of two hydrogen atoms and an oxygen atom. As mentioned earlier, when two or more elements combine, the resulting molecule is referred to as a compound. image

Compounds that contain the elements carbon and hydrogen are classified as organic, and all others as inorganic. Living tissues are based on organic compounds, but the body requires inorganic compounds too.

Covalent and ionic bonds

The vast array of chemical processes on which life is based is completely dependent upon the way atoms come together, bind and break apart. For example, the simple water molecule is a crucial foundation of all life on Earth. If water was a less stable compound, and the atoms came apart easily, human biology could never have evolved. On the other hand, the body is dependent upon the breaking down of various molecules (e.g. sugars, fats) to release energy for cellular activities. When atoms are joined together, they form a chemical bond that is generally one of two types: covalent or ionic.

Covalent bonds are formed when atoms share their electrons with each other. Most molecules are held together with this type of bond; it forms a strong and stable link between its constituent atoms. A water molecule is built using covalent bonds. Hydrogen has one electron in its outer shell, but the optimum number for this shell is two. Oxygen has six electrons in its outer shell, but the optimum number for this shell is eight. Therefore, if one oxygen atom and two hydrogen atoms combine, each hydrogen atom will share its electron with the oxygen atom, giving the oxygen atom a total of eight outer electrons, making it stable. The oxygen atom shares one of its electrons with each of the two hydrogen atoms, so that each hydrogen atom has two electrons in its outer shell, and they too are stable (Fig. 2.4).

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Figure 2.4 A water molecule, showing the covalent bonds between hydrogen (yellow) and oxygen (green).

Ionic bonds are weaker than covalent bonds and are formed when electrons are transferred from one atom to another. For example, when sodium (Na) combines with chlorine (Cl) to form sodium chloride (NaCl) there is a transfer of the only electron in the outer shell of the sodium atom to the outer shell of the chlorine atom (Fig. 2.5).

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Figure 2.5 Formation of the ionic compound, sodium chloride.

This leaves the sodium atom of the compound with eight electrons in its outer (second) shell, and therefore stable. The chlorine atom also has eight electrons in its outer shell, which, although not filling the shell, is a stable number. The sodium atom is now positively charged because it has given away a negatively charged electron, and the chloride ion is now negatively charged because it has accepted sodium’s extra electron. The two atoms, therefore, stick together because they are carrying opposite, mutually attractive, charges.

When sodium chloride is dissolved in water the ionic bond breaks and the two atoms separate. The atoms are charged, because they have traded electrons, so are no longer called atoms, but ions. Sodium, with the positive charge, is a cation, written Na+, and chloride, being negatively charged, is an anion, written Cl. By convention the number of electrical charges carried by an ion is indicated by the superscript plus or minus signs. image

Electrolytes

An ionic compound, e.g. sodium chloride, dissolved in water is called an electrolyte because it conducts electricity. Electrolytes are important body constituents because they:

• conduct electricity, essential for muscle and nerve function

• exert osmotic pressure, keeping body fluids in their own compartments

• act as buffers to resist pH changes in body fluids.

Many biological compounds, e.g. carbohydrates, are not ionic, and therefore have no electrical properties when dissolved in water. Important electrolytes other than sodium and chloride include potassium (K+), calcium (Ca2+), bicarbonate (HCO3) and phosphate (PO42−).

Molecular weight

The molecular weight of a molecule is the sum of the atomic weights of the elements forming its molecules, e.g.:

Water (H2O)

2 hydrogen atoms

(atomic weight 1)

2

1 oxygen atom

(atomic weight 16)

16

 

Molecular weight

= 18

Sodium bicarbonate (NaHCO3)

1 sodium atom

(atomic weight 23)

23

1 hydrogen atom

(atomic weight 1)

1

1 carbon atom

(atomic weight 12)

12

3 oxygen atoms

(atomic weight 16)

48

 

Molecular weight

= 84

Molecular weight, like atomic weight, is expressed simply as a figure until a scale of measurement of weight is applied.

Molarity

This is the commonest way to express the concentration of many substances in body fluids.

The mole (mol) is the molecular weight in grams of a substance. One mole of any substance contains 6.023 × 1023 molecules or atoms. For example, 1 mole of sodium bicarbonate (the example above) is 84 g.

In a molar solution, 1 mole of a substance is dissolved in 1 litre of solvent (dissolving fluid). In the human body the solvent is usually water. A molar solution of sodium bicarbonate is therefore prepared by dissolving 84 g of sodium bicarbonate in 1 litre of solvent.

Molar concentration may be used to measure quantities of electrolytes, non-electrolytes, ions and atoms provided the molecular weight of the substance is known. It means that a molar solution of a substance contains exactly the same number of particles as any other molar solution. If the molecular weight of the substance is not known, or if there is more than one material in solution, another system of measuring concentration has to be used, such as grams per litre. The tiny quantities of many substances dissolved in body fluids mean that physiological concentrations are often expressed as fractions of a mole: millimoles/litre (thousandths of a mole) or micromoles/litre (millionths of a mole) (Table 2.2). Table 2.3 gives examples of the normal plasma levels of some important substances, given in molar concentrations and alternative units.

Table 2.2 Molar concentrations

Solute units

Quantity per litre of solvent

1 mole of sodium chloride molecules (NaCl)

58.5 g

1 millimole of sodium chloride molecules

0.0585 g (58.5 mg)

1 mole of sodium ions

23 g

1 micromole of sodium ions

0.000023 g (23 μg)

1 mole of carbon atoms

12 g

1 mole of oxygen gas (O2)

32 g

Table 2.3 Examples of normal plasma levels

Substance

Amount in moles

Amount in other units

Chloride

97–106 mmol/l

97–106 mEq/l*

Sodium

135–143 mmol/l

135–143 mEq/l*

Glucose

3.5–5.5 mmol/l

60–100 mg/100 ml

Iron

14–35 mmol/l

90–196 mg/100 ml

Acids, alkalis and pH

The concentration of hydrogen ions ([H+]) in a solution is a measure of the acidity of the solution. Control of hydrogen ion levels in body fluids is an important factor in maintaining a stable internal environment.

An acid substance releases hydrogen ions when in solution. On the other hand, a basic (alkaline) substance accepts hydrogen ions, often with the release of hydroxyl (OH) ions. A salt releases other anions and cations when dissolved; sodium chloride is therefore a salt because in solution it releases sodium and chloride ions.

* Milliequivalents per litre (mEq/l)

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Concentration is expressed as:

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The pH scale

The standard scale for measurement of hydrogen ion concentration in solution is the pH scale. The scale measures from 0 to 14, with 7, the midpoint, as neutral; this is the pH of water. Water is a neutral molecule, neither acid nor alkaline, because when the molecule breaks up into its constituent ions, it releases one H+ and one OH, which balance one another. Most body fluids are close to neutral, because strong acids and bases are damaging to living tissues, and body fluids contain buffers, themselves weak acids and bases, to keep their pH within narrow ranges.

A pH reading below 7 indicates an acid solution, while readings above 7 indicate alkalinity (Fig. 2.6). A change of one whole number on the pH scale indicates a 10-fold change in [H+]. Therefore, a solution of pH 5 contains ten times as many hydrogen ions as a solution of pH 6.

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Figure 2.6 The pH scale.

Not all acids ionise completely when dissolved in water. The hydrogen ion concentration is, therefore, a measure of the amount of dissociated acid (ionised acid) rather than of the total amount of acid present. Strong acids dissociate more freely than weak acids, e.g. hydrochloric acid dissociates freely into H+ and Cl, while carbonic acid dissociates much less freely into H+ and HCO3.

Likewise, not all bases dissociate completely. Strong bases dissociate more fully, i.e. release more OH than weaker ones.

pH values of the body fluids

The pH of body fluids must be maintained within relatively narrow limits depending on the fluid concerned. The normal range of pH values of some body fluids are shown in Table 2.4.

Table 2.4 pH values of certain body fluids

Body fluid

pH

Blood

7.35 to 7.45

Saliva

5.4 to 7.5

Gastric juice

1.5 to 3.5

Bile

6 to 8.5

Urine

4.5 to 8.0

The highly acid pH of the gastric juice is maintained by hydrochloric acid secreted by the parietal cells in the walls of the gastric glands. The low pH of the stomach fluids destroys microbes and toxins that may be swallowed in food or drink. Saliva has a pH of between 5.4 and 7.5, which is the optimum value for the action of salivary amylase, the enzyme present in saliva which initiates the digestion of carbohydrates. Amylase is destroyed by gastric acid when it reaches the stomach.

Blood pH is kept between 7.35 and 7.45, and outwith this narrow range there is severe disruption of normal physiological and biochemical processes. Normal metabolic activity of body cells produces certain acids and alkalis, which would tend to alter the pH of the tissue fluid and blood. Chemical buffers are responsible for keeping body pH stable.

Buffers

Despite the constant cellular production of acid and alkaline substances, body pH is kept stable by systems of buffering chemicals in body fluids and tissues. These buffering mechanisms temporarily neutralise fluctuations in pH, but can function effectively only if there is some means by which excess acid or alkali can be excreted from the body. The organs most active in this way are the lungs and the kidneys. The lungs are important regulators of blood pH because they excrete carbon dioxide (CO2). CO2 increases [H+] in body fluids because it combines with water to form carbonic acid, which then dissociates into a bicarbonate ion and a hydrogen ion.

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The lungs, therefore, help to control blood pH by regulating levels of excreted CO2. The brain detects rising [H+] in the blood and stimulates breathing, causing increased CO2 loss and a fall in [H+]. Conversely, if blood pH becomes too alkaline, the brain can reduce the respiration rate to increase CO2 levels and increase [H+], restoring pH towards normal (see Ch. 10).

The kidneys regulate blood pH by increasing or decreasing the excretion of hydrogen and bicarbonate ions as required. If pH falls, hydrogen ion excretion is increased and bicarbonate conserved; the reverse happens if pH rises. In addition, the kidneys generate bicarbonate ions as a by-product of amino acid breakdown in the renal tubules; this process also generates ammonium ions, which are rapidly excreted.

Other buffer systems include body proteins, which absorb excess H+, and phosphate, which is particularly important in controlling pH inside cells. The buffer and excretory systems of the body together maintain the acid–base balance so that the pH range of the blood remains within normal, but narrow, limits.

Acidosis and alkalosis

The buffer systems described above compensate for most pH fluctuations, but these reserves are limited and, in extreme cases, can become exhausted. When the pH falls below 7.35, and all the reserves of alkaline buffers are used up, the condition of acidosis exists. In the reverse situation, when the pH rises above 7.45, the increased alkali uses up all the acid reserve and the state of alkalosis exists.

Acidosis and alkalosis are both dangerous, particularly to the central nervous system and the cardiovascular system. In practice, acidotic conditions are commoner than alkalotic ones, because the body tends to produce more acid than alkali. Acidosis may follow respiratory problems, if the lungs are not excreting CO2 as efficiently as normal, or if the body is producing excess acids (e.g. diabetic ketoacidosis, p. 228) or in kidney disease, if renal H+excretion is reduced. Alkalosis may be caused by loss of acidic substances through vomiting, diarrhoea, endocrine disorders or diuretic therapy, which stimulates increased renal excretion. Rarely, it may follow increased respiratory effort, such as in an acute anxiety attack where excessive amounts of CO2 are lost through overbreathing (hyperventilation).

Important biological molecules

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Learning outcomes

After studying this section, you should be able to:

image describe in simple terms the chemical nature of sugars, proteins, lipids, nucleotides and enzymes

image discuss the biological importance of each of these important groups of molecules.

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Carbohydrates

Carbohydrates (sugars and starches) are composed of carbon, oxygen and hydrogen. The carbon atoms are normally arranged in a ring, with the oxygen and hydrogen atoms linked to them. The structures of glucose, fructose and sucrose are shown in Figure 2.7. When two sugars combine to form a bigger sugar, a water molecule is expelled and the bond formed is called a glycosidic linkage.

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Figure 2.7 The combination of glucose and fructose to make sucrose.

Glucose, the main form in which sugar is used by cells, is a monosaccharide. Monosaccharides can be linked together to form bigger sugars, ranging in size from two sugar units (disaccharides), e.g. sucrose (Fig. 2.7) or table sugar, to long chains containing many thousands of monosaccharides e.g. starch. Such complex carbohydrates are called polysaccharides.

Glucose can be broken down (metabolised) in either the presence (aerobically) or the absence (anaerobically) of oxygen, but the process is much more efficient when O2 is used. During this process, energy, water and carbon dioxide are released (p. 307). To ensure a constant supply of glucose for cellular metabolism, blood glucose levels are tightly controlled. The sugars:

• provide a ready source of energy to fuel cell metabolism (p. 307)

• provide a form of energy storage, e.g. glycogen (p. 307)

• form an integral part of the structure of DNA and RNA (pp. 429431)

• can act as receptors on the cell surface, allowing the cell to recognise other molecules and cells.

Amino acids and proteins

Amino acids always contain carbon, hydrogen, oxygen and nitrogen, and many in addition carry sulphur. In human biochemistry, 20 amino acids are used as the principal building blocks of protein, although there are others; for instance, there are some amino acids used only in certain proteins, and some are seen only in microbial products. Of the amino acids used in human protein synthesis, there is a basic common structure, including an amino group (NH2), a carboxyl group (COOH) and a hydrogen atom. What makes one amino acid different from the next is a variable side chain. The basic structure and three common amino acids are shown in Figure 2.8. As in formation of glycosidic linkages, when two amino acids join up the reaction expels a molecule of water and the resulting bond is called a peptide bond.

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Figure 2.8 Amino acid structures: A. Common structure, R = variable side chain. B. Glycine, the simplest amino acid. C. Alanine. D. Phenylalanine.

Proteins are made from amino acids joined together, and are the main family of molecules from which the human body is built. Protein molecules vary enormously in size, shape, chemical constituents and function. Many important groups of biologically active substances are proteins, e.g.:

• carrier molecules, e.g. haemoglobin (p. 59)

• enzymes (p. 24)

• many hormones, e.g. insulin (p. 218)

• antibodies (pp. 371–372).

Proteins can also be used as an alternative energy source, usually in starvation. In starvation, the main source of body protein is muscle tissue, so it is accompanied by wasting of muscles.

Lipids

The lipids are a diverse group of substances whose common property is an inability to mix with water (i.e. they are hydrophobic). They are made up of carbon, hydrogen and oxygen atoms. The most important groups of lipids include:

• phospholipids, integral to cell membrane structure. They form a double layer, providing a water-repellant barrier separating the cell contents from its environment (p. 28)

• certain vitamins (p. 270). The fat-soluble vitamins are A, D, E and K

• fats (triglycerides), stored in adipose tissue (p. 37) as an energy source. Fat also insulates the body and protects internal organs. A molecule of fat contains three fatty acids attached to a molecule of glycerol (Fig. 2.9). When fat is broken down under optimal conditions, more energy is released than when glucose is fully broken down.

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Figure 2.9 Structure of a fat (triglyceride) molecule.

Fats are classified as saturated or unsaturated, depending on the chemical nature of the fatty acids present. Saturated fat tends to be solid, whereas unsaturated fats are fluid.

• prostaglandins are important chemicals derived from fatty acids and are involved in inflammation (p. 367) and other processes.

• cholesterol is a lipid made in the liver and available in the diet (p. 269). It is an integral part of cell membranes and is used to make steroid hormones (p. 216).

Nucleotides

Nucleic acids

These are the largest molecules in the body and are built from nucleotides. They include deoxyribonucleic acid (DNA, p. 429) and ribonucleic acid (RNA, p. 431).

Adenosine triphosphate (ATP)

ATP is a nucleotide that contains ribose (the sugar unit), adenine (the base) and three phosphate groups attached to the ribose (Fig. 2.10A). It is sometimes known as the energy currency of the body, which implies that the body has to ‘earn’ (synthesise) it before it can ‘spend’ it. Many of the body’s huge number of reactions release energy, e.g. the breakdown of sugars in the presence of O2. The body captures the energy released by these reactions, using it to make ATP from adenosine diphosphate (ADP). When the body needs chemical energy to fuel cellular activities, ATP releases its stored energy, water and a phosphate group through the splitting of a high-energy phosphate bond, and reverts to ADP (Fig. 2.10B).

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Figure 2.10 ATP and ADP: A. Structures. B. Conversion cycle.

The body needs chemical energy to:

• drive synthetic reactions (i.e. building biological molecules)

• fuel movement (locomotion)

• transport substances across membranes.

Enzymes

Many of the body’s chemical reactions can be reproduced in a test-tube. Surprisingly, the rate at which the reactions then occur usually plummets to the extent that, for all practical purposes, chemical activity ceases. The cells of the body have developed a solution to this apparent problem – they are equipped with a huge array of enzymes. Enzymes are proteins that act as catalysts for biochemical reactions – that is, they speed the reaction up but are not themselves changed by it, and therefore can be used over and over again. Enzymes are very selective and will usually catalyse only one specific reaction. The molecule(s) entering the reaction is called the substrate and it binds to a very specific site on the enzyme, called the active site. Whilst the substrate(s) is bound to the active site the reaction proceeds, and once it is complete the product(s) of the reaction breaks away from the enzyme and the active site is ready for use again (Fig. 2.11).

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Figure 2.11 Action of an enzyme: A. Enzyme and substrates. B. Enzyme–substrate complex. C. Enzyme and product.

Enzyme action is reduced or stopped altogether if conditions are unsuitable. Increased or decreased temperature is likely to reduce activity, as is any change in pH. Some enzymes require the presence of a cofactor, an ion or small molecule that allows the enzyme to bind its substrate(s). Some vitamins are cofactors in enzyme reactions.

Enzymes can catalyse both synthetic and breakdown reactions, and their names (almost always!) end in ˜ase. When an enzyme catalyses the combination of two or more substrates into a larger product, this is called an anabolic reactionCatabolic reactions involve the breakdown of the substrate into smaller products, as occurs during the digestion of foods.

Movement of substances within body fluids

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Learning outcomes

After studying this section, you should be able to:

image compare and contrast the processes of osmosis and diffusion

image using these concepts, describe how molecules move within and between body compartments.

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Movement of substances within and between body fluids, sometimes across a barrier such as the cell membrane, is essential in normal physiology.

From a physical point of view, substances will always travel from an area of high concentration to one of low concentration, assuming that there is no barrier in the way. Between two such areas, there exists a concentration gradient and movement of substances occurs down the concentration gradient, or downhill, until concentrations on each side are equal (equilibrium is reached). No energy is required for such movement, so this process is described as passive.

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There are many examples in the body of substances moving uphill, i.e. against the concentration gradient; in this case, chemical energy is required, usually in the form of ATP. These processes are described as active. Movement of substances across cell membranes by active transport is described on page 32.

Passive movement of substances in the body proceeds usually in one of two main ways – diffusion or osmosisimage

Diffusion

Diffusion refers to the movement of a chemical substance from an area of high concentration to an area of low concentration, and occurs mainly in gases, liquids and solutions. Sugar molecules heaped at the bottom of a cup of coffee that has not been stirred will, in time, become evenly distributed throughout the liquid by diffusion (Fig. 2.12). The process of diffusion is speeded up if the temperature rises and/or the concentration of the diffusing substance is increased.

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Figure 2.12 The process of diffusion: a spoonful of sugar in a cup of coffee.

Diffusion can also occur across a semipermeable membrane, such as the plasma membrane or the capillary wall. Only molecules able to cross the membrane will be able to diffuse through. For example, oxygen diffuses freely through the walls of the alveoli (airsacs in the lungs), where oxygen concentrations are high, into the bloodstream, where oxygen concentrations are low. However, blood cells and large protein molecules in the plasma are too large to cross and so remain in the blood.

Osmosis

While diffusion of solute molecules across a semipermeable membrane results in equal concentrations of the solute on both sides of the membrane, osmosis refers specifically to diffusion of water down its concentration gradient. This is usually because the solute molecules are too large to pass through the pores in the membrane. The force with which this occurs is called the osmotic pressure. Imagine two solutions of sugar separated by a semipermeable membrane whose pores are too small to let the sugar molecules through. On one side, the sugar solution is twice as concentrated as on the other. After a period of time, the concentration of sugar molecules will have equalised on both sides of the membrane, not because sugar molecules have diffused across the membrane, but because osmotic pressure across the membrane pulls water from the dilute solution into the concentrated solution – i.e. water has moved down its concentration gradient. Osmosis proceeds until equilibrium is reached, at which point the solutions on each side of the membrane are of the same concentration and are said to be isotonic. The importance of careful control of solute concentrations in the body fluids can be illustrated by looking at what happens to a cell (e.g. a red blood cell) when it is exposed to solutions that differ from normal physiological conditions.

Plasma osmolarity is maintained within a very narrow range because if the plasma water concentration rises, i.e. the plasma becomes more dilute than the intracellular fluid within the red blood cells, then water will move down its concentration gradient across their membranes and into the red blood cells. This may cause the red blood cells to swell and burst. In this situation, the plasma is said to be hypotonic. Conversely, if the plasma water concentration falls so that the plasma becomes more concentrated than the intracellular fluid within the red blood cells (the plasma becomes hypertonic), water passively moves by osmosis from the blood cells into the plasma and the blood cells shrink (Fig. 2.13).

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Figure 2.13 The process of osmosis. Net water movement when a red blood cell is suspended in solutions of varying concentrations (tonicity): A. Isotonic solution. B. Hypotonic solution. C. Hypertonic solution.

Body fluids

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Learning outcomes

After studying this section, you should be able to:

image define the terms intra- and extracellular fluid

image using examples, explain why homeostatic control of the composition of these fluids is vital to body function.

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The total body water in adults of average build is about 60% of body weight. This proportion is higher in babies and young people and in adults below average weight. It is lower in the elderly and in obesity in all age groups. About 22% of body weight is extracellular water and about 38% is intracellular water (Fig. 2.14).

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Figure 2.14 Distribution of body water in a 70 kg person.

Extracellular fluid

The extracellular fluid (ECF) consists mainly of blood, plasma, lymph, cerebrospinal fluid and fluid in the interstitial spaces of the body. Other extracellular fluids are present in very small amounts; their role is mainly in lubrication, and they include joint (synovial) fluid, pericardial fluid (around the heart) and pleural fluid (around the lungs).

Interstitial or intercellular fluid (tissue fluid) bathes all the cells of the body except the outer layers of skin. It is the medium through which substances pass from blood to the body cells, and from the cells to blood. Every body cell in contact with the ECF is directly dependent upon the composition of that fluid for its well-being. Even slight changes can cause permanent damage, and ECF composition is, therefore, closely regulated. For example, a fall in plasma potassium levels may cause muscle weakness and cardiac arrhythmia, because of increased excitability of muscle and nervous tissue. Rising blood potassium also interferes with cardiac function, and can even cause the heart to stop beating. Potassium levels in the blood are only one of the many parameters under constant, careful adjustment by the homeostatic mechanisms of the body.

Intracellular fluid

The composition of intracellular fluid (ICF) is largely controlled by the cell itself, because there are selective uptake and discharge mechanisms present in the cell membrane. In some respects, the composition of ICF is very different from ECF. Thus, sodium levels are nearly ten times higher in the ECF than in the ICF. This concentration difference occurs because, although sodium diffuses into the cell down its concentration gradient, there is a pump in the membrane that selectively pumps it back out again. This concentration gradient is essential for the function of excitable cells (mainly nerve and muscle). Conversely, many substances are found inside the cell in significantly higher amounts than outside, e.g. ATP, protein and potassium.

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