unit 3 understanding chemistry

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Unit 3: Understanding matter - atoms and molecules

Unit 3: Understanding matter – atoms and molecules Page 3-2

Case study .........................................................................................................................................4

Objectives..........................................................................................................................................5

What is matter? .................................................................................................................................6

The Particle theory .................................................................................................................................... 6

States of matter ........................................................................................................................................ 7

Chemical and physical changes ................................................................................................................. 9

Classification of matter ........................................................................................................................... 11 Homogeneous matter ......................................................................................................................... 11 Heterogeneous mixtures .................................................................................................................... 11 Special types of mixtures .................................................................................................................... 12

Building blocks of matter ................................................................................................................. 14

Elements .................................................................................................................................................. 14

Atoms ...................................................................................................................................................... 15

Isotopes ................................................................................................................................................... 17

Electron orbitals - energy levels and valence electrons .......................................................................... 18

Ions .......................................................................................................................................................... 19

Molecules ................................................................................................................................................ 20

Bonding ........................................................................................................................................... 22

Ionic bonding ........................................................................................................................................... 22

Covalent bonding .................................................................................................................................... 23

Rules to predict which type of bonding will occur ................................................................................... 24

Multiple bonds ........................................................................................................................................ 25

Cyclic bonding ......................................................................................................................................... 26

Aromatic bonding ................................................................................................................................... 27

Hydrogen bonding ................................................................................................................................... 27

Unit 3: Understanding matter – atoms and molecules -3

Chemical reactions ........................................................................................................................... 29

Energy of reactions ................................................................................................................................. 30 Catalysts .............................................................................................................................................. 32 Enzymes – catalysts in the body ......................................................................................................... 33

Measuring atoms and molecules – the mole ..................................................................................... 34

Mass – mole conversions. ................................................................................................................ 35

Solutions ......................................................................................................................................... 37

Dilutions .................................................................................................................................................. 40 Calculating dilution factors ................................................................................................................. 41 Preparing diluted solutions ................................................................................................................. 42

Electrolytes ...................................................................................................................................... 43

Acids and Bases ............................................................................................................................... 45

What is meant by pH? ............................................................................................................................. 47

Acid-base neutralisation ......................................................................................................................... 48

Buffers ..................................................................................................................................................... 48

Organic compounds ......................................................................................................................... 50

Biomolecules ........................................................................................................................................... 51 Carbohydrates ..................................................................................................................................... 54 Lipids ................................................................................................................................................... 54 Amino acids and proteins ................................................................................................................... 55

Bibliography .................................................................................................................................... 56


Unit 3: Understanding matter - atoms and molecules

Case study Fiona presented at outpatients with symptoms of extreme lethargy, rapid breathing and abdominal pain. After an initial consultation it was revealed that Fiona suffers from insulin dependent diabetes mellitus and had missed taking her last insulin dose. Blood tests show a pH of 7.32 along with high concentrations of the electrolyte potassium and elevated glucose, ketones and lactic acid levels. Urine pH was very low at 4.3. These results along with an electrocardiograph showed erratic heart muscle activity confirmed that Fiona was in a state of metabolic acidosis and may be at risk of going into a coma.

Immediate treatment for Fiona was administration of 10 mL of 10% calcium gluconate to reduce cardiac muscle excitability along with insulin and 50 mL 50% glucose. To help her breathing, Fiona was given ventolin.

Once stabilised Fiona will need to take regular doses of insulin and maintain a healthy diet comprising of approximately 50% carbohydrates, 15% protein, less than 10% unsaturated fats, limited sugar and salt intake. Cholesterol should be maintained at or below 5.0 mmol/L and blood glucose (after fasting) to be 4.7 mmol/L.

Understanding Fiona’s case: The chemistry involved in Fiona’s case requires an understanding of concentrations, acids, bases and buffers, chemical reactions, enzymes and types of biomolecules such as proteins.

It is hoped that through studying this unit you will appreciate that the human body is made up of an array of molecules and that the body functions through a series of chemical reactions. When the body is unable to maintain a healthy state, treatment often includes the administration of medications at appropriate doses.


Objectives On completion of this section you should be able to:

• name and describe the states of matter

• use the particle theory to discuss matter

• distinguish between chemical and physical changes

• classify matter into pure substances and mixtures

• describe an atom and its structure using terms such as nucleus, proton, neutron and electron

• explain the meaning of the terms atomic number, mass number, isotopes and atomic weight

• discuss why atoms bond to form molecules

• explain what is meant by a molecular formula

• explain aspects of chemical reactions such as reversibility, heat of reaction, role of catalysts and enzymes

• define the terms solution, solute, solvent and solubility

• describe what is meant by the term concentration and explain its importance in relation to the reporting of test results from clinical pathology laboratories

• perform dilution calculations using a simple formula

• define the terms electrolyte and non-electrolyte

• define the terms acid, base and buffer

• describe the pH scale of acidity

• explain the difference between organic and inorganic compounds

• name some biomolecules present in the body

.


What is matter? Matter is defined as being anything that occupies space and has mass. Mass is a measure of the amount of matter that an object contains. For example, a golf ball has more mass than a table tennis ball, and therefore contains more matter.

But what does matter consist of?

Before answering this question you are invited to perform the following mini-practical and reflect on possible explanations for the observations that you make.

Mini-practical 3-1: Diffusion - Perfumes, dyes in water

Equipment: You will need some food colouring and a glass of water

1. Recall how perfumes, nasty smells and nice food odours spread quickly in air.

2. Place a couple of drops of food colouring in a glass of water and observe how the dye moves.

3. Can you explain your observations? … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … …… … …

… … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … …… … …

… … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … …… … …

What is happening? The spread of odours and colours suggests outward movement of particles from a concentrated source.

Consider what happens when you mix a bucket of sand and a bucket of gravel. The combined mix does not equate to 2 buckets. On a micro scale this phenomenon can be observed using methanol and water. Mixing ½ a cup of water with ½ a cup of methanol results in a total volume of less than 1 cup. This observation is best explained by assuming that particles of water and methanol fit between each other.

The particle theory The observations made in these mini-practicals suggest very strongly that matter is composed of particles. The particle theory is an accepted scientific theory that is used to explain many phenomena.

The particle theory states that:

• Matter is composed of particles.

• Particles of matter are in constant random motion.

• The speed at which particles move increases with increasing temperatures.

• Particles of matter are held together by electrostatic forces.

• Empty spaces exist between the particles of matter.


All matter is composed of very small particles, called atoms (or groups of atoms called molecules). The distance particles are apart and the forces of attraction between the particles determine the form or state that the matter exists in.

States of matter Matter exists in three different physical states. These physical states of matter are: solid, liquid and gas. If the particles of matter are strongly attracted to each other and held close together, it is in the solid state. When the forces of attraction are reduced but particles are still held closely together, matter is in the liquid state. If the distance between the particles is large and forces of attraction are negligible matter takes a gaseous form.

Consider what happens when a corn kernel is heated in oil and forms popcorn. As the kernel is heated, the particles of moisture inside gain energy and the liquid changes to gas. The pressure inside the kernel increases to a point where the kernel ruptures. The starch and proteins in the kernel also expand and the texture becomes like airy foam. Refer to figure 3-1.

Figure 3-1: Corn kernel and popped popcorn

Although there are three physical states of matter, it is important to realise that chemical substances are unchanged in these physical states. The only difference between them is in how closely they are packed together (the density of the substance which is the ratio of mass to volume) and in how rapidly these molecules are moving (their kinetic energy).

Table 3-1 summarises the important characteristics of the three states of matter.


Table 3-1: The three states of matter—solids, liquids and gases

State of matter Characteristics Visual representation

Solid Matter that has a fixed shape and volume (e.g. coal, sugar, bone, iron)

When matter is in the solid state, the particles are held together closely in fixed positions by attractive forces. There is very little movement of these particles. Since the particles are already packed together, a solid is not easily compressed.

Solid

Liquid Matter that has a fixed volume and takes the shape of a container (e.g., milk, blood, water)

Particles in a liquid although held closely together are able to move past each other, allowing the liquid to flow. Like solids, liquids cannot be easily compressed. Liquid

Gas Matter that takes both the shape and the volume of its container.

In the gas state particles of matter are at a great distance apart and have negligible forces of attraction between each other. This allows gases to expand without limit to fill any space and also explains why gases are easily compressed. (e.g., oxygen)

Gas

(Adapted from Hickman & Caon, 1995.)

Matter can be transformed from one state to another by adding or removing energy in the form of heat. When a solid becomes a liquid it is said to melt and the temperature at which this happens is the melting point. When a liquid reaches boiling point, it transforms to a gas. This process is known as evaporation. As heat is removed, a gas turns to liquid and is said to condense. As more heat is removed and the liquid turns to a solid we say it freezes.

Consider the process of evaporation. As heat is added to a liquid, the particles or molecules gain more energy and are able to move faster and further apart until they are able to overcome the forces of attraction that hold them together. When this occurs, the state of matter is that of a gas.


The processes involved in change of state of matter are summarised in Figure 3-2.

Melting Evaporation

Freezing Condensation

SOLID LIQUID GAS

Figure 3-2: Processes of change of state of matter

Chemical and physical changes The change that occurs as matter changes state between solid, liquid and gas are physical changes. In physical changes no new substance is formed since the substance is the same, it just has different physical properties.

A chemical change occurs when new substances are formed that have different properties and composition to the original material.

So as to witness chemical and physical changes first hand you might like to perform the following tasks. Identify each task as being an example of either a physical or chemical change and suggest on a particle level what may be occurring.

Mini-practical 3-2: Chemical and physical changes

Equipment: You will need an egg, vinegar, a thermometer and a glass of hot water.

1. Place an egg into a container of vinegar and leave for 48 hours. Observe any changes that may occur to the appearance and feel of the shell.

… … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … …… … …

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What is happening? The composition of the egg shell has changed chemically. The calcium carbonate in the egg shell has reacted with the acetic acid in the vinegar. This chemical reaction results in the removal of calcium from the egg shell making it go soft, or even disappear.


2. Place a (non digital) thermometer into a glass of hot water and observe any changes that are seen in the thermometer’s mercury or alcohol level over 2 hours.

… … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … …… … …

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… … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … …… … …

What is happening? The mercury/alcohol level changes as a result of the temperature change of the water. When the water is hot, the particles of mercury/alcohol have greater energy and are able to move faster and at a greater distance apart. As the water cools, mercury/alcohol particles have less energy and so move closer together. The physical change in the mercury/alcohol allows it to be used effectively as an indicator of temperature in the thermometer.

Activity 3-1

Indicate in the following table whether the occurrences are physical or chemical changes.

Occurrence Physical or chemical change?

Explanation

Melting ice

Burning paper

Rusting iron

Tearing paper

Breaking glass

Boiling water

Bleaching a stain

Fermenting grapes


Classification of matter Some materials are composed of numerous types of matter (variable composition), while others consist of a single type of matter (constant composition).

Homogeneous matter When matter has a constant or uniform composition, it is homogeneous. Homogenous matter has only one phase, each portion of the sample being identical in composition, resulting in the sample having the same composition throughout. Homogeneous matter can be either a homogeneous mixture (such as a solution where the solid is completely dissolved in the liquid as in a teaspoon of sugar in a cup of tea) or a pure substance (such as gold).

Some other examples of homogenous mixtures include:

• Salt water – a single phase, uniform mixture of salt dissolved in water.

• A piece of steel – a single phase, uniform mixture of different metals forming an alloy.

• A solution of bleach - a single phase, uniform mixture of hydrogen peroxide in water.

Homogeneous mixtures can be separated by physical processes into pure substances.

Pure substances are composed of elements or compounds. The smallest unit or particle of matter is the atom. Elements contain only one type of atom while compounds contain two or more types of atoms in fixed proportions. Pure substances cannot be broken down into smaller units by physical methods. Compounds can only be separated into elements by chemical methods.

Heterogeneous mixtures When matter has variable composition it is a heterogeneous mixture. The components of a heterogeneous mixture are easily separated and have different properties. A heterogeneous mixture is composed of mixed or multiple phases. Heterogeneous mixtures can be separated by physical processes into homogeneous matter.

Fruit salad is a good example of a heterogeneous mixture – it is composed of many different and distinguishable solid phases. Fruit salad can be separated into apples, grapes and strawberries etc. by hand.

Some other examples of heterogeneous mixtures include:

• The sand at the beach – composed of different solid phases as many different coloured sand particles can be observed.

• Soft drink – composed of a gas and liquid phase as bubbles of carbon dioxide are present in the drink.

• Blood – composed of different liquid phases since the components of blood are red and white blood cells, platelets and plasma which are easily separated when left to stand or when centrifuged (spun at high speed).


Activity 3-2

Classify the following as either homogenous or heterogeneous mixtures:

Mixture Homogeneous or heterogeneous?

Explanation

Jelly

Air when clouds are present

Air in a room

Salad dressing of vinegar and oil

Special types of mixtures Colloids and emulsions are names given to particular types of homogeneous mixtures. Remember a homogeneous mixture appears to have the same composition throughout since the particle sizes are too small to be seen. Colloids are mixtures where one substance is evenly dispersed throughout another. There are many examples, including gas/liquid colloids such as fog, or aerosol sprays; solid/gas colloids such as whipped cream or liquid/liquid colloids such as hand cream or homogenised milk. Liquid/liquid colloids are given a special name of emulsion.

A suspension is a particular type of heterogeneous mixture, since it has large particles that settle out on standing. An example of a suspension is a preparation of the antibiotic, amoxicillin or augmentin, in water where the active ingredient is mixed with water but does not dissolve, rather remains as a suspended solid: this is why it is important to shake these medicines thoroughly before measuring a dose.

Figure 3-3 provides a summary of information regarding the classification of matter.


Figure 3-3: Classification of matter

Separated physically

Separated physically

Separated chemically

Matter (solids, liquids and gases)

Heterogeneous mixture (variable composition)

Homogeneous matter (uniform composition)

Pure substance (only one substance)

Homogeneous mixture (two or more substances of

uniform composition)

Element (one type of atom)

Compound (two or more types of atoms in

fixed proportions)

Atom


Building blocks of matter

Elements As mentioned previously, the elements are pure substances containing only one type of atom. Many atoms of the element need to be grouped together before macroscopic properties of the element are evident (e.g. lustre of metals). Elements cannot be broken down into simpler substances by chemical methods. There are approximately 92 naturally occurring elements and a number of synthetic ones. Many elements will be familiar to you, for example gold, silver, iron and carbon. All of the elements are listed systematically in the Periodic Table. A copy of the Periodic Table of the Elements is provided in Figure 3-4.

Figure 3-4: The Periodic Table of the elements (Adapted from Timberlake 1992)

Consider the layout of the periodic table. The horizontal rows of the periodic table are called periods (or rows). There are seven periods. And the vertical columns are called groups. There are 18 groups.

The Periodic Table is designed so as to group elements with similar properties together. This grouping system simplifies a study of the elements since knowledge of a particular group rather than every element (118 of them!) is usually all that is needed to predict the chemistry of the element. For example the chemical behaviour of the elements in group 1A (the alkali metals) with the exception of hydrogen, is very similar. Other groups that should be noted are group 2A (the alkaline earth metals), group7A (the halogens) and group 8A the noble gases. Note also that there is a zig zag black line going from the top left to bottom right of the periodic table. Elements on the left of this line (except for hydrogen) are metals, while those to the right of the line are non-metals. With the


exception of aluminium, those elements located beside the stepped line are metalloids, which have some properties of both metals and nonmetals.

You will notice that the names for the elements listed in the Periodic Table have been abbreviated using a form of shorthand notation. Elements are abbreviated using any of the following:

• the first letter (e.g. hydrogen = H),

• the first two letters (e.g. helium = He)

• the first letter and another letter of the element name (e.g. chlorine = Cl) or

• letters based on the Latin names of the element (e.g. potassium = K from the latin, Kalium)

Table 3-2, provided for your interest only, presents the elements that are essential for the body to function.

Table 3-2: Essential elements for the body

Name of element Symbol Elements necessary in trace amounts

Symbol

Carbon C Chlorine Cl

Hydrogen H Cobalt Co

Oxygen O Fluorine F

Nitrogen N Iodine I

Iron Fe Magnesium Mg

Calcium Ca zinc Zn

Phosphorus P Copper Cu

Potassium K Selenium Se

Sodium Na Manganese Mn

Sulphur S Molybdenum Mo

Adapted from Strobe, 2008.

Atoms The smallest unit of an element is the atom. If an atom is split (as in a nuclear reactor) that atom no longer exists – different atoms are formed. The atom is made up of a central, extremely dense nucleus which is surrounded by a cloud of rapidly moving, extremely small electrons. The nucleus contains two types of particles, the proton and the neutron. (A representation of an atom is given in Figure 3-5). The electrons move around the nucleus at a larger distance and at high speed creating an electron cloud. This cloud can be likened to the appearance of the spinning blades of a helicopter. When the rotor is spinning, the two blades appear as if they occupy a circular cloud around the rotor. Electrons move in paths called electron energy shells or electron orbitals.


Figure 3-5: The arrangement of the subatomic particles in an atom. All of the subatomic particles are extremely small compared to the things you see around you. Because the mass of a subatomic particle is so minute, chemists find it convenient to use a very small unit of mass called an atomic mass unit (amu). A single proton and a single neutron each have a mass of 1 amu. Since the electron is so small and light, its mass is usually ignored in atomic mass calculations.

The electron is much lighter than either the proton or neutron (approximately 2000 electrons equal the weight of one proton). Table 3-3, summarises some of the information known about the subatomic particles. (Note the use of scientific notation to write the mass of each particle.)

The proton is positively charged, the electron, negatively charged and the neutron has no charge. All atoms are electrically neutral as they contain the same number of protons as electrons.

Table 3-3: Characteristics of the subatomic particles in atoms

Subatomic particle

Electrical charge (relative values)

Particle mass Atomic mass units (amu) Location in atom

Proton +1 1.67263 × 10-24 g 1.00728 Inside nucleus

Neutron 0 1.67494 × 10-24 g 1.00867 Inside nucleus

Electron –1 9.10939 × 10-28 g 5.48580 × 10-4 Outside nucleus

Adapted from Gilbert, Kirss & Davies , 2004.

Atoms of each element are identified by the number of protons present in the nucleus. The number of protons in a nucleus of an atom is defined by the atomic number. Different atoms have different atomic numbers. For example any atom with 1 proton will be a hydrogen atom, any atom with 2 protons will be a helium atom and any atom with 6 protons will be a carbon atom. Since atoms are neutral, the atomic number is also equivalent to the number of electrons.

Another number which identifies an atom is the mass number. The mass number indicates the mass of the atom. Since the mass is due to protons and neutrons, it is possible to calculate the number of neutrons present in an atom by subtracting the atomic number (number of protons) from the mass number (number of protons and neutrons).


Consider some examples of using atomic and mass numbers to determine the number of protons, neutrons and electrons in an atom:

• Hydrogen has an atomic number of 1. This means that an atom of hydrogen has 1 proton and 1 electron. The mass number of hydrogen is also 1 which indicates that hydrogen does not have any neutrons.

• Helium has an atomic number of 2, and therefore each helium atom contains 2 protons and 2 electrons. Since helium has a mass number of 4, there must also be 2 neutrons present.

• An atom of potassium with an atomic number of 19 and mass number of 39 has 19 protons and 19 electrons (inferred from the atomic number) and 20 neutrons, (calculated by subtracting atomic number from mass number, i.e. 39 – 19=20).

The Periodic Table is arranged so that the atomic numbers increase from left to right across the table, and from top to bottom down the table.

Activity 3-3

Complete the following table:

Atom Atomic number

Mass number Number of protons

Number of electrons

Number of neutrons

Carbon 6 6

Sulfur 16 16

Fluorine 9 19

Oxygen 16 8

Sodium 11 23

Isotopes Atoms of the same element always have the same number of protons present in the nucleus, however not all atoms of the same element have the same number of neutrons.

Isotopes are atoms of the same element that have different numbers of neutrons. The atomic numbers of the isotopes are the same, but their mass numbers are different. When we are talking about isotopes we say the name of the element and then the mass number, e.g. carbon 13. When we write isotopes, the shorthand convention is to place the mass number as a superscript in front of the symbol, e.g. 13C. or it can also be written 13

6C, with both mass number and atomic numbers indicated.

Most of the elements occur in nature as mixtures of isotopes. However, many isotopes are unstable. Indeed, some are radioactive and are continuously decomposing to form other elements. Of the seven known isotopes of carbon, only two, carbon-12 and carbon-13 are stable. You may have heard of carbon dating, which analyses the amount of carbon-14 present in carbon containing specimens. Other isotopes are used in radiation therapy. Unit 5 will introduce applications of radiation in medicine.


Electron orbitals - energy levels and valence electrons The chemical behaviour of an element is determined by its electronic configuration – i.e. how electrons are arranged around the nucleus of the atom.

The cloud of electrons that surrounds the nucleus has electrons arranged in a series of seven discrete energy levels or orbitals. Orbitals closest to the nucleus have the lowest energy. Each energy level holds a different number of electrons. The maximum number of electrons in the first three energy levels is shown in the Table 3-4.

Table 3-4: Electrons in first three energy levels

Electron energy level Maximum number of electrons 1 2 2 8 3 18

For the first 18 elements in the Periodic Table, energy levels tend to be filled with electrons consecutively. This means that 2 electrons enter the first energy level before any electrons enter the second energy level. Similarly, 8 electrons enter the second energy level before electrons begin to enter the third energy level. Note that so far a total of 10 electrons have been accounted for, so the next 8 electrons will enter the3rd energy level.

The electrons in the outer energy level are referred to as valence electrons. These valence electrons have an important role in determining the chemical properties of the element.

Activity 3-4

a) Using the information from the Periodic Table (Atomic number) and the maximum number of electrons allowed in the first 3 energy levels, complete the following table by predicting how many valence electrons the following elements have.

Element Atomic number Number of electrons Number of valence electrons

Fluorine

Carbon

Oxygen

Chlorine

Hydrogen

Neon

Sodium

Sulfur


b) Of the elements in this table, there are three pairs, each pair belonging to a different group of the Periodic Table. Can you identify the three pairs and name the group that each pair belongs to?

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c) Do you notice a trend with the number of valence electrons for elements in the same group of the Periodic Table?

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Ions Atoms may lose or gain valence electrons in order to attain the electronic configuration of a noble gas. If electrons are gained, the atom becomes negatively charged, and if electrons are lost, the atom becomes positively charged. These charged species are called ions. Positive ions are referred to as cations while negative ions are referred to as anions.

Whether an atom forms an anion or cation is dependant on their relevant position in the Periodic Table. Metals (elements positioned on the left hand side of the stepped line) tend to lose electrons and become positively charged ions or cations. Non-metals (those elements positioned on the right hand side of the stepped line) tend to gain electrons and become negatively charged ions or anions. Following these guidelines you will notice that the gain or loss of electrons is kept at a minimum.

Activity 3-5

Using the position in the Periodic Table and the number of valence electrons for each of the elements in Activity 3-4, predict what type of ion will form and the number of charges on it.

Element

Position in Periodic Table

(metal/non-metal?)

Number of valence

electrons

Anion or cation?

Number of charges on the

ion

Fluorine

Oxygen

Chlorine

Hydrogen

Neon

Sodium

Sulfur


Molecules Atoms do not generally occur singly but combine, in fixed proportions, with other atoms to form a molecule.

The molecule exists as a single unit with its own unique set of characteristics, different to the combining atoms. Remember in the flowchart describing the breakdown of matter, compounds (which are molecules) were listed as pure substances that contain two or more types of atoms in fixed proportions.

Molecular formulas are as a shorthand method to represent the atoms present and the ratio in which they combine to form the molecule. Some molecules that you will be familiar with are water and oxygen.

Water forms when 2 atoms of hydrogen combine with one atom of oxygen. Since water is a combination of hydrogen and oxygen in fixed ratio of 2:1, the molecular formula for water is H2O, the subscript 2 indicating that 2 hydrogen atoms are present.

The oxygen that we breathe is molecular oxygen O2, (not atomic oxygen!). The molecular formula for oxygen indicates that 2 atoms of oxygen combine to form the molecule.

Figure 3-6 provides details of how numbers are used in the molecular formula.

Indicates theelement sodium

(one atom)

Indicates theelement chlorine

(one atom)

NaC

Indicates theelement

hydrogen

Indicates twoatoms ofhydrogen

Indicates theelement sulfur

(one atom)

Indicates fouratoms ofoxygen

Indicates theelementoxygen

SO2 4H

Indicates theelement calcium

(one atom)

Indicates the nitrategroup composed of

one nitrogen atom andthree oxygen atoms

Indicates twonitrate groupsCa(NO 3 2)

Figure 3-6: Molecular formulas


The names and molecular formulas for other molecules that you may encounter are shown in Table 3-5 and Activity 3-6:

Table 3-5: Ratio of combining atoms for some molecules

Molecule Molecular formula Ratio of combining atoms

Sodium chloride NaCl 1atom of sodium : 1 atom of chlorine

Sulfuric acid H2SO4 2 atoms of hydrogen : 1 atom of sulfur : 4 atoms of oxygen

Calcium nitrate Ca(NO3)2 1 calcium atom : 2 nitrogen atoms : 6 oxygen atoms

Glucose C6H12O6 6 atoms of carbon : 12 atoms of hydrogen : 6 atoms of oxygen

Sometimes large molecules are formed when a molecule has other molecules associated with it. For example calcium sulfate dihydrate is a calium sulfate molecule that has 2 molecules of water associated with it. It is represented as CaSO4.2H2O. Note the use of the dot which implies that there are two molecules of water associated with every molecule of calcium sulphate in the crystal structure of the molecule. (These water molecules are termed “water of crystallisation”. )

Activity 3-6

Count the number of atoms in each molecular formula given in the following table

Molecule Molecular formula Ratio of combining atoms

Vitamin C C6H8O6

Heparin C12H19NO20S3

Aspirin C9H8O4

Ventolin (C13H21NO3)2.H2SO4


Bonding Molecules form when two or more atoms bond chemically so that the product formed is more stable or at a lower energy than the combining atoms. Since elements with a full outer energy level are very stable (the noble gases), other atoms will also become very stable if they can obtain a filled outer energy level. Atoms can gain a full outer energy level by bonding with other atoms.

Bonds are formed between atoms through electrostatic interactions. A chemical bond forms when particular atoms share, gain or lose valence electrons. Bonding occurs because the resulting molecule is of lower energy than the combining atoms.

Since the Noble gases already have a full outer energy level they are not involved in bonding.

When studying bonding in molecules, it is generally only necessary to consider the valence electrons of combining atoms.

Ionic bonding When electrons are donated from one atom or group of atoms to another an ionic bond forms.

Sodium chloride, NaCl, is formed when sodium donates an electron and chlorine accepts an electron, see Figure 3-7. Look back at your answers to Activity 3-5. Sodium had 1 valence electron, and tended to lose this electron and form a cation Na+, so as to achieve noble gas configuration, while chlorine had 7 valence electrons, and tended to gain an electron and form an anion Cl-, in order to attain noble gas configuration. Sodium donates its electron to chlorine and forms an ionic bond.

(Note: the dots in the following diagrams represent electrons in the valence energy level – these dot diagrams are called Lewis diagrams or Lewis structures)

Figure 3-7: Ionic bonding in sodium chloride


Covalent bonding When a group of atoms share the electrons around their nuclei, a covalent bond is formed.

Hydrogen gas is formed when two hydrogen atoms share their electrons. In Activity 3-5, you may have predicted that hydrogen has 1 valence electron, and needs one more electron in order to attain noble gas configuration. When two hydrogen atoms each share their 1 valence electron, each hydrogen atom attains noble gas configuration and a covalent bond is formed. The covalent bond consists of a pair of electrons being shared between two atoms. The single covalent bond can be represented by a single line drawn between the two hydrogen atoms, as shown in Figure 3-8.

Alternative representation

Figure 3-8: Covalent bonding in hydrogen When electrons are shared between the same atoms, such as in hydrogen, H2, they are shared equally. This type of bond is a pure covalent bond. However when electron sharing occurs between different types of atoms, such as in water, H2O, the sharing is not equal and the bond is referred to as a polar covalent bond.

Consider the formation of covalent bonds in methane as shown in Figure 3-9. Methane has the molecule that consists of one carbon atom and hydrogen atoms. Since carbon only has 4 valence electrons, it will need a share in another 4 electrons in order to gain a full valence shell (i.e. a valence shell of 8 electrons). Since a hydrogen atom has only 1 valence electron, it is easy to predict that there must be 4 hydrogen atoms that bond to carbon, to form methane.

Alternative representation

Figure 3-9: Covalent bonding in methane Consider the formation of covalent bonds in water as shown in Figure 3-10 Oxygen has 6 valence electrons (solid dots) while each hydrogen has only 1 electron (open dots). The two covalent bonds are formed when oxygen share 1 electron from each of the hydrogen atoms. In this way, oxygen has a share in a total of 8 electrons and gains a full valence shell, and each hydrogen atom has a share in 2 electrons and also gains a full valence shell.


Figure 3-10: Covalent bonding in water In the case of water, the oxygen atom has a greater share in the electrons, so the oxygen end of the molecule has a partial negative charge (designated δ–) and the hydrogen end of the molecule has a partial positive charge (designated δ+).

Water is a polar molecule and is often said to be a polar solvent. The polar nature of water is illustrated in Figure 3-11.

O

H H

δ+δ+

δ- δ-

Figure 3-11: Covalent bonding in water

Rules to predict which type of bonding will occur There are three simple rules that can be used to predict the type of bonding that is likely to occur between elements, namely:

1. Noble gas elements do not bond at all, e.g. Helium

2. Metals bond with non metals by ionic bonds, e.g. Na and Cl combine to form salt sodium chloride.

3. Elements that are non-metals, from the same groups or groups that are close together form covalent bonds, e.g. H and O combine to form water. The further the groups are apart on the Periodic Table the more polar the bond will be.


Activity 3-7

Predict what type of bonding will occur between atoms of

a) hydrogen and sulfur… … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … …

b) two oxygen atoms… … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … …

c) beryllium and chlorine… … … … … … … … … … … … … … … … … … … … … … … … … … … … … … …

d) argon… … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … …

Multiple bonds In the discussion of covalent bonding, a bond forms between two atoms when a pair of electrons is shared. Sharing of one pair of electrons forms a single covalent bond. Sometimes it is necessary for two atoms to share two or even three pairs of electrons. If this occurs double or triple covalent bonding occurs. Carbon dioxide, (CO2), is a molecule that has double bonds between the central carbon atom and each oxygen atom. The carbon atom has 4 valence electrons and each oxygen atom has 6 valence electrons.

If two pairs of electrons were shared between each oxygen atom and the central carbon atom, all atoms would have a complete set of valence electrons. When two pairs of electrons are shared a double bond is formed. The double bond can be represented by two parallel lines drawn between the bonding atoms, as shown in Figure 3-12.

Figure 3-12: Covalent bonding in carbon dioxide (formation of double bonds) A triple bond forms when three pairs of electrons are shared. The triple bond can be represented by three parallel lines drawn between the bonding atoms. Activity 3-8 deals with triple bonds.


Activity 3-8

Acetylene has the molecular formula C2H2. Use diagrams to represent the valence electrons for carbon and hydrogen atoms and show how a triple bond forms between the two carbon atoms in acetylene.

The chemical structure of many drugs, such as steroids, involves cyclic and aromatic bonding.

Cyclic bonding Some molecules, such as cyclohexane, form by bonding in ring structures – think of it as one end of the molecule bonding to the other. Ring structures can be symbolised as in figure 3-13. The ring represents the carbon backbone of the molecule. Since carbon has 4 valence electrons, it will need a share in another 4 electrons to attain a full valence shell. Each carbon will form 4 bonds. In the ring structures, each corner represents a carbon atom, and since each carbon bonds to two other carbons, there will also be two hydrogen atoms attached to each carbon.

CC

C

CC

C H

H HH

HHH

HH

H

H

H

Figure 3-13: Representations of cyclic covalent bonds in cyclohexane, C6H12

(2 dimensional representation and simple ring representation)


Aromatic bonding A special type of ring structure can also form which includes double bonds, alternating around the ring. In benzene, each carbon is involved in a total of 3 bonds with other carbons, and therefore bonds to only one hydrogen atom. This type of ring structure is termed an aromatic ring, and is depicted in Figure 3-14. Note that in the aromatic ring structures, each corner represents a carbon atom and the one hydrogen atom attached to it.

H H

H

HH

H

Figure 3-14: Representations of aromatic bonds in benzene, C6H6.

Hydrogen bonding Hydrogen bonding is a special type of attractive force that forms between molecules.

Recall that the water molecule is polar, with the oxygen slightly negative and the hydrogen slightly positive. A molecule of water can position itself towards the partial charges of other water molecules so as to maximise any electrical attractions. The electrical attractions that occur between the molecules of water are called hydrogen bonds. Figure 3-15 shows how hydrogen bonds form between the oxygen atom (slightly negative) of one water molecule, and a hydrogen atom (slightly positive) of other water molecules.

Figure 3-15: Hydrogen bonds (the dotted lines) occur between the oxygen atom of one water molecule and the hydrogen atom of other water molecules


Although they are referred to as hydrogen ‘bonds’, the attraction between the partial charges is much much weaker that the attraction between atoms involved in ionic or covalent bonds.

Hydrogen bonding explains why ice is able to float on water. In liquid water, hydrogen bonds form between water molecules bringing the molecules closer together. In ice, molecules have less energy and movement. Water molecules in ice are aligned in a lattice structure and are not able to move to allow hydrogen bonding to occur. Consequently hydrogen bonds do not form to the same extent, as they do in liquid water. This makes water denser than ice and explains why ice floats on water.

Activity 3-9

Can you explain why a full bottle of water expands to the point of bursting when it is frozen?

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Chemical reactions When bonds between atoms or molecules are formed or broken, a chemical reaction is said to have occurred. Previously in this unit, we looked at physical and chemical changes. For every chemical change that occurs, or whenever a new substance if formed, a chemical reaction has taken place. An example of a chemical change is that of iron rusting. The chemical reaction that is occurring in this situation is the interaction of iron, oxygen and water to form iron oxide.

Chemists use chemical equations, which are a shorthand notation, to record a great deal of information about the reaction.

The chemical equation for the reaction of iron rusting can be written as:

( ) ( ) ( ) ( )2 2 3 2solid 2 gas liquid solid4Fe +3O +6H O 2Fe O .3(H O)→ Eqn 3-1

The chemical equation conveys a great deal of information about the reaction that is occurring, in regard to the reactants (ingredients) required to form the product.

On an atomic or molecular level the reactants are:

• 4 Fe = 4 atoms of solid iron

• 3O2 = 3 molecules of oxygen gas (oxygen gas is O2) and

• 6 H2O = 6 molecules of liquid water

These reactants combine to form the product:

• 2 Fe2O3.3(H2O) = 2 molecules of ferric oxide (hydrated)

Each molecule of ferrous oxide consists of 2 atoms of iron, 3 atoms of oxygen with 3 molecules of water attached.

Additional information is given by the

• type of arrow used for the reaction:

The use of a single arrow ‘→ ’ indicates that the reaction will only proceed in the direction written. Such a reaction is irreversible.

The use of a double arrow ‘ ’, indicates that the reaction is reversible and will occur in both directions.

• Indication of the physical state of the reactants and products:

The physical state of the reactants and products are indicated in brackets after the formulae in the equation – usually abbreviations of s, l, g and aq are used for solid, liquid, gas and aqueous respectively. (Note the term aqueous is used to indicate a solution where water is the solvent.)


The reaction of salt water with calcium carbonate to form sodium carbonate crystals and calcium chloride is an example of a reversible reaction.

( ) 3( ) 2 3( ) 2( )2NaCl +CaCO Na CO +CaClaq s s aq Eqn 3-2

It is important to realise that these equations are balanced in terms of atoms reacting and atoms in the product. This is a consequence of the Law of Conservation of Mass which states that matter cannot be created or destroyed.

We can use the equation to say that 2 molecules of NaCl react with 1 molecule of CaCO3 to produce 1 molecule each of Na2CO3 and CaCl2 . And, keeping the same molecular ratios 4 molecules of NaCl react with 2 molecules of CaCO3 to produce 2 molecules each of Na2CO3 and CaCl2 .

Activity 3-10

Verify that the equations used as examples are balanced by counting the number of atoms present.

Reactants Product Number of atoms of:

( ) ( ) ( )2solid 2 gas liquid4Fe +3O +6H O ( )2 3 2 solid2Fe O .3(H O)

Fe

O

H

Reactants Products Number of atoms of: ( ) 3( )2NaCl +CaCOaq s 2 3( ) 2Na CO +CaCls (aq)

Na

Cl

Ca

C

O

Energy of reactions (endothermic and exothermic reactions; activation energy)

When chemical reactions occur, energy is either absorbed from the surroundings or released to the surroundings through the breaking and formation of bonds. This energy or heat of reaction can be used to classify reactions as exothermic or endothermic reactions.

If you would like to see an exothermic reaction, try mini-practical 3-3.


Mini-practical 3-3: Exothermic reactions

Equipment: You will need a jar with a lid, a thermometer, steel wool and vinegar.

1. Find a jar with a lid, so that the lid can be put on when a thermometer is left inside the jar.

2. Place a thermometer in the jar and record the temperature after 5 minutes.

3. Soak some steel wool in vinegar for a few minutes.

4. Squeeze as much vinegar as possible out of the steel wool

5. Wrap the steel wool around the bulb of the thermometer and place back into the jar. Replace the lid.

6. Read the thermometer after 5 or 10 minutes and note any temperature changes that occur.

7. If a temperature change has occurred, remove the steel wool and observe any change to its appearance.

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What is happening? Vinegar is used to remove any protective coating from the steel wool, allowing the steel to rust in a moist environment. As the reaction proceeds a temperature rise is observed. This indicates that heat is released from the reaction.

When heat is released to the surroundings the reaction is said to be exothermic. The word exothermic can be broken down to “exo” meaning out, and “thermic” meaning heat. The rusting of iron was exothermic as was shown by the increase in temperature in the steel wool mini-practical. Other examples of exothermic reactions are the combustion (burning) of fossil fuels such as coal, petroleum and natural gas. If a reaction produces a lot of heat very quickly, then an explosion may occur. In exothermic reactions, the products have lower energy than the reactants.

Other chemical reactions can involve the absorption of energy, more specifically the absorption of light energy, from the surroundings. An example of this type of reaction is the production of glucose by photosynthesis. Reactions in which heat is absorbed are endothermic reactions. Again, the word can be broken down to “endo” meaning in and “thermic” meaning heat. In endothermic reactions, the products have higher energy than the reactants. The net energy difference between the breaking of bonds in reactants and formation of bonds for the product determines whether a reaction will be endothermic or exothermic.

Most chemical reactions do not take place instantaneously. There is an energy barrier (a ‘hill’), called the activation energy, that must be overcome before molecules can react. Molecules will only react if they collide with sufficient energy to overcome this barrier. The energy needed comes from the kinetic energy of the molecules. Kinetic energy increases with increasing temperature so chemical reactions speed up as the temperature of the reactants increases.


Catalysts Another method used to increase the speed of a reaction is to lower the activation energy, making it more likely for molecules to react. The activation energy of a reaction can be lowered by using a catalyst. A catalyst works by providing an alternative reaction pathway of lower energy. Note that a catalyst does not change the amount or type of products, only the speed at which they are produced. A catalyst is a substance that increases the rate (or speed) of a chemical reaction without itself being consumed. Figures 3-16a and 3-16b, illustrate the activation energy of a reaction and how a catalyst reduces the energy required for the reaction to occur. The use of a catalyst for a chemical reaction can be likened to the alternate route of a steam train that allows it to go through a tunnel through a mountain, rather than climb the mountain in order to get to the same destination. Less energy is required if the tunnel is used, and it is more likely that the destination could be reached. The use of a catalyst lowers the activation energy making it easier for the reaction to proceed to form products.

Figure 3-16 a): Use of a tunnel as the alternate route, making it easier for the steam train to reach its destination.

Figure 3-16b): Dashed line shows effect of a catalyst in lowering the activation energy of a reaction


Enzymes – catalysts in the body The cells in the human body contain special catalysts called enzymes. Nearly all the chemical reactions that occur in living tissue require an enzyme to proceed to an acceptable metabolic rate. The enzyme works by bringing reactants together, making it easier for the reaction to occur and form products.

The action of an enzyme is illustrated in Figure 3-17. Note that the enzyme is unchanged in the process.

Figure 3-17: How an enzyme works

Pineapple contains an enzyme that breaks down protein. You may like to see the action of this enzyme by conducting Mini-practical 3-4 which investigates the effect of this enzyme on the setting ability of jelly – gelatine in jelly is a protein.

Mini-practical 3-4: Enzyme action

Equipment: You will need fresh pineapple, tinned pineapple, packet jelly, 4 bowls.

1. Label the bowls, fresh, tinned, boiled and control

2. Boil some fresh pineapple pieces and allow to cool

3. Prepare jelly as per instructions on packet.

4. Place 6 pieces of each type of pineapple into the correspondingly marked bowl – no pineapple in the control.

5. Pour hot jelly solution into each bowl, ensuring that the pineapple is covered.

6. Place bowls in fridge and allow jelly to set.

7. What conclusions can you draw from this experiment? … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … … …… … …

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What is happening? The jelly in the control bowl sets as expected since no enzyme is present to interfere with the action of the protein in gelatine. The jelly in the bowls with boiled and tinned pineapple also sets as the boiling and canning procedures on the pineapple break down the enzyme. The jelly with fresh pineapple will not set as the enzyme has broken down the protein in the gelatine.

Measuring atoms and molecules – the mole In our day to day lives we keep track of the quantity of different items by mass, (e.g. a kilogram of meat), volume (e.g. litres of petrol) or number (a dozen eggs).

In society, we use a number of different counting units when referring to groups of items,

A pair of socks = 2 socks A dozen eggs = 12 eggs A ream of paper = 500 sheets A grand = $1000

You may notice that the number each term refers to is relevant to the useable quantity of that item. Chemists refer to the mass of 1 proton or 1 neutron as 1 atomic mass unit (amu). Using relative atomic masses, in which the simplest atom hydrogen has a relative atomic mass of 1 amu, it follows that oxygen has a relative atomic mass of approximately 16 amu and carbon approximately 12 amu.

The size or mass of an atom is so small that it is impossible to work with (e.g. the mass of hydrogen atom is 1.67 x 10-24 g).

(1.67 x 10-24 g = or 1.67 x 10-27 kg = 0.000 000 000 000 000 000 000 000 001 67 kg! This is an extremely small mass!)

Dealing with such small masses, made it necessary for chemists to devise some form of counting atoms or molecules that was relevant to their size. The unit that is used to do this is the ‘mole’.

A mole of an element has the same mass in grams as the atomic mass of that element. This mass is termed the molar mass. Since hydrogen has an atomic mass of 1, 1 mole of hydrogen atoms has a molar mass of 1 g; carbon with an atomic mass of 12, will have a molar mass of 12 g.

A mole of a molecule has the same mass in grams as the sum of the atomic masses of the elements that form the compound. This mass is termed the molar or molecular mass of the molecule. The molar mass of a molecule of water (H2O) is equivalent to

2 x atomic mass of H + atomic mass of O = 2 x 1 + 16 = 18 g

International standards apply to atomic masses with the standard carbon-12 being assigned an atomic mass of exactly 12 grams.


With the most accurate measurements available today, it has been determined that there are

6.022 ×1023 atoms of carbon-12 in 12.0000 g of pure carbon-12. A mole is equivalent to 6.022 x 1023 units. The number 6.022 × 1023 is called Avogadro’s number (NA) after a famous nineteenth century chemist.

To appreciate the size of the mole, consider the following:

• A mole of marshmallows would cover the earth to a depth of 1 km.

• Counting at 1 object per second, it would take 19 trillion years to count to a mole.

• 1 mole of dollar coins distributed to everyone on earth – would be enough to spend $1 million per hour 24 hours a day and still have half left over at end of lifetime!!

Activity 3-11

Use the information provided in the Periodic Table to determine

a) The mass of 1 mole of sulfur atoms … … … …… … …

b) The mass of 1 mole of carbon atoms … … … …… … …

c) The molar mass of a molecule of hydrochloric acid (HCl) … … … …… … …

d) The molar mass of a molecule of water (H2O) … … … …… … …

e) The molar mass of a molecule of glucose (C6H12O6) … … … …… … …

Mass – mole conversions.

If the mass of an element or compound is known, the number of mole can be calculated using the appropriate atomic or molar mass. In chemistry, the number of mole of a substance is referred to as the amount of the substance.

1

mass of substance(g)amount of substance, n (no. mole) = molar mass (g mole )− eqn 3-3

Consider the following examples: 1. How many mole of carbon is present in 20g?

Solution:

1

20 g carbonamount of substance, n (no. mole) = 12 g mole

1.67mole of carbon

−

=


2. How many mole of iron oxide(Fe2O3) is present in 0.5 kg?

Solution:

Since the units of molar mass are gram per mole it is necessary to convert the mass of 0.5 kilogram to gram. Note 0.5 kg = 500 g

Molar mass of Fe2O3 = 2 x 55.85gmol-1 (due to Fe) + 3 x 16 g mole-1 (due to O)=159.7 g mole-1

2 31

500 g Fe Oamount of substance, n (no. mole) = 159.7 g mole

3.13 mole of iron oxide

−

=

3. How many gram of sodium chloride is present in 0.50 mole?

Solution:

The molar mass of NaCl = 22.99 + 35.45 = 58.44g mole-1. Note that equation 3-3 can be rearranged as

1mass of substance(g) molar mass (g mole ) amount of substance, n (no. mole)x−=

Mass NaCl (g) = 58.44g mole-1 x 0.50 mole = 29.22 g

Activity 3-12

Complete the following table, of mole to mass conversions.

Substance Molar mass (g mol-1) Amount (mole) Mass (g)

Mg (magnesium) 24.31 0.46

H2O (water) 500

C6H12O6 (glucose) 180.156 225

CH3OH (methanol) 0.25


Solutions In the discussion regarding classification of matter, we introduced the terms homogeneous mixture and heterogeneous mixture, to describe solutions, colloids, emulsions and suspensions. In almost all applications it is important to have an idea of how much of one substance is mixed into another.

Of particular importance is the concentration or strength of solutions. Remember that a solution results when one substance is dissolved in another. Components of the solution are the solute and solvent. You would be familiar with the concentrated juice that can be purchased in supermarkets. This juice is too concentrated to drink as is but is suitable after it is diluted with water. (Usually 500 mL of concentrated juice is diluted with 1500 mL of water to make a total volume of 2 L.)

A solution is a type of homogenous mixture in which one substance called the solute is dispersed uniformly in another substance called the solvent. The solubility of a solute in a solvent is an indication of the maximum amount of solute that will completely dissolve in a given volume of the solvent. A saturated solution results when this limit of solubility is reached. The rate at which a solute dissolves in a solvent can be increased by stirring or applying heat and by ensuring the particle size of the solute is small (all of which allows greater mixing and exposure of solute particle in the solvent). If you would like to investigate some aspects of solubility, perform Mini-practical 3-5.


Mini-practical 3-5: Saturated solutions

Equipment: You will need cold water, saucepan, sugar and access to a hot plate.

1. Measure 1 cup of cold water into a saucepan and add sugar, a teaspoon at a time, until the sugar no longer dissolves.

2. How many teaspoons were required? … … … … … … … … … … … … … … … … … … … … … …

3. Place the saucepan on a hot plate and heat gently, stirring the solution.

4. Has the sugar dissolved? … … … … … … … … … … … … … … … … … … … … … …

5. What conclusions can you draw from this experiment?

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What is happening? There is a limit to the number of teaspoons of sugar that will dissolve in a cup of cold water. Heating the saturated solution gives the particles more energy, increasing movement and allowing more sugar to dissolve.

If one substance or solute is soluble in another this means that it will dissolve in the solvent. Some vitamins are water-soluble, e.g. vitamin C and the B group vitamins. This means that they dissolve in water and are generally not stored in the body. Fat-soluble vitamins such as vitamin D and E dissolve in body fat rather than water and are retained in the body.

Concentration of solutions (i.e. the amount of solute per solvent) can be expressed in a variety of units, including:

• mass per volume = mass of solute per unit volume of solvent. e.g. g/L, mg/mL • %w/v = mass of solute (in gram) per 100 mL of solvent • %w/w = mass of solute (in gram) per 100 g of solvent • %v/v = volume of solute (in mL) per 100 mL solvent • molarity = moles of solute per litre of solvent (mol/L, mol L-1, or M). For very dilute

solutions the terms millimolar or micromolar may be used. Remember milli and micro mean 10-3 and 10-6 respectively.

• ppm = parts per million. This generally refers to a ratio by mass, i.e. 1 mass unit ( gram) of solute per million mass units (grams) of solvent, but since water is usually the solvent and has a density of 1 g per mL, 1ppm = 1mg per litre. (i.e. 1 mg per million mg, since 1 litre of water has a mass of 1kg = 1 000 000 mg).

• ppb = parts per billion. Similar to ppm except it refers to gram of solute per billion gram of solvent which is equivalent to microgram solute per litre solvent. (i.e. 1 microgram per billion micrograms, since 1 litre of water has a mass of 1 000 000 000 µg).

• molality = moles solute per kilogram of solvent


It may be easier to relate to the units of concentration with a few examples, as shown in Table 3-7.

Table 3-7: Common concentration units and examples

Concentration unit Example Explanation

Mass per volume 125 mg /5 mL amoxicillin 125 mg amoxicillin per 5 mL of solution

%w/v normal saline (NS) solution of 0.9% w/v NaCl

0.9 grams per 100 mL of solution

%w/w Medicated acne gel of 2.5% w/w benzoyl peroxide

2 .5 grams of benzoyl peroxide in every 100 gram of gel material

%v/v 70% v/v ethanol solution 70 mL of ethanol per 100 mL of total volume of solution

molar, M 1 M glucose solution 1 mole glucose per litre of solution (check 1 mole glucose C6H12O6 = 180 g glucose) To prepare such a solution you would dissolve:

180 g of glucose in 1 litre of water (or

90g of glucose in 500 mL of water or

18 g of glucose in 100 mL of water.)

millimolar, mM 200 mM zinc gluconate solution

200 milli moles or 0.2 moles of zinc gluconate per 1 litre of water.

micromolar, µM 1 µM dose of DLV (delavirdine mesylate)

A dose with the equivalent of 1 micromole or 0.000 001 mole of DLV per litre or water.

ppm Acceptable level of lead in drinking water is <0.015ppm

Less than 0.015 mg lead per litre of water

(1 litre water has a mass of 1 million milligrams)

ppb Acceptable level of arsenic in drinking water is < 10 ppb

Less than 10 microgram per litre of water

(1 litre water has a mass of 1 billion micrograms)


Dilutions

When a volume of a solution is diluted, the concentration is decreased. This is logical since in order to dilute something, more solvent must be added to it.

The new concentration of the solution can be calculated by considering the dilution factor. If the original solution was diluted by a factor of ½, the new concentration would be ½ the original. This would be a 2-fold dilution. Likewise a 10-fold dilution or 1/10 dilution would result in a solution with a concentration 1/10 of the original. These dilutions are illustrated in Figure 3-18.

Figure 3-18: 2-fold and 10-fold dilutions

Mini-practical 3-6: Dilutions

Equipment: You will need cups, cordial, water and measuring cups.

Prepare the following series of cordial drinks from a concentrated cordial solution – make sure that each diluted solution is mixed well.

1. ½ (or 2-fold)

2. 1/3 (or 3-fold)

3. 1/5 (or 5-fold)

4. 1/10 (or 10-fold)

5. Comment on how the strength of these solutions is indicated by their colour?

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What is happening? The darker the colour of the solution, the more concentrated it is.


Calculating dilution factors Sometimes it will be necessary to calculate how to dilute a solution to a specific concentration. The dilution factor may be calculated by expressing the two concentrations as a fraction:

new concentrationdilution factor =original concentration

Eqn 3-4a

(and of course express this fraction in lowest terms!) .

Or by considering the change in volume that occurs.

initial solution volumedilution factor =final solution volume

Eqn 3-4b

Consider the dilutions shown in Figure 3-18. In the first example, the original concentration was 4 mg/mL and the diluted concentration was 2 mg/mL. The dilution factor is calculated as:

=new concentrationdilution factor =

original concentration2 14 2=

In the second example the original concentration was 50 mg/mL and the diluted concentration was 5 mg/mL. The dilution factor is calculated as:


original concentration5 1

50 10=

In preparing these solutions, specially calibrated volumetric glassware is used so as to ensure accurate measurements. The volume of concentrated solution is measured using a pipette before it is delivered into a volumetric flask. The flask has a mark on it which indicates the level to which water should be added to achieve the desired concentration of dilute solution.

Activity 3-13

Calculate the dilution factor required when a stock saline solution of 2.0% v/v is to be diluted to 0.5 % v/v?

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… … … … … … … … … … … … … … … … … … … … … … … … … … … … … …


Preparing diluted solutions In diluting solutions you will always have a stock solution (the original most concentrated solution) and the required concentration for the dilute solution that you are to prepare.

If the stock solution has a concentration of 100 mg/L and you need to prepare a solution that is 50 mg/L, It should be obvious that this new diluted solution is ½ the strength of the original. So the dilution factor is ½.

The dilution factor can then be used to determine how to prepare any particular volume of the diluted solution.

For example if we wish to prepare 20ml of the dilute solution, ½ of 20mL = 10mL, so we use 10 mL of stock solution and dilute it (or make it up) to 20mL with water.

If we wish to prepare 70mL of dilute solution, ½ of 70mL = 35mL, so use 35mL of stock and dilute up to 70mL.

Consider another example.

If a stock solution has a concentration of 12.5% w/v and we need to prepare 60mL of a 3%w/v solution how do we do it? Firstly calculate the dilution factor,


original concentration3 0.24

12.5=

Next use the dilution factor to calculate how much stock solution is required in the dilution.

volume of stock solution required = dilution factor × required volume of dilute solution

In this case, volume of stock solution required = 0.24 X 60mL = 14.4 mL

This means that we take 14.4 mL of the stock solution and make up to 60 mL. The new solution will have a concentration of 3% w/v and will be a total of 60mL.

To summarise these steps…

1. Calculate the dilution factor ..

new concentrationdilution factor =original concentration

2. Use the dilution factor to determine how much of the original stock solution is required to make the required volume of dilute solution.

volume of stock solution required = dilution factor × required volume of dilute solution


Electrolytes Chemical substances that conduct an electric current when dissolved in water are called electrolytes. Electrolytes have the ability to conduct electricity because they carry an electrical charge and have the ability to move freely through a solution. Examples of electrolytes include ionic compounds, salts, acids, and bases.

Note that many chemicals don’t conduct electricity when dissolved in water; these chemicals are called non-electrolytes. For example, sugar (sucrose) is a non-electrolyte.

Solid sodium chloride is an ionic compound because it is formed when sodium loses an electron and chlorine gains an electron. The solid is then held together by electrostatic forces between the positive sodium ions and the negative chloride ions. When sodium chloride (table salt) dissolves in water it seems to disappear. This is because the solid separates into its constituent ions and become surrounded by water molecules, as shown in Figure 3-19.

Figure 3-19: Salt (NaCl) dissolving in water

The chemical equation that describes what is happening is:

NaCl(s) + H2O(l) → Na+(aq) + Cl-

(aq) Eqn 3-5

The ions make it possible for a solution of sodium chloride to conduct electricity.

Your body fluids -- blood, plasma, interstitial fluid (fluid between cells) -- are rather like seawater and have a high concentration of sodium chloride (NaCl). The electrolytes in sodium chloride are:

• sodium ion (Na+) – a cation

• chloride ion (Cl-) – an anion


Most of the body’s electrolytes are dissolved in the body fluids. The remainder are found attached to proteins or deposited as solids in the form of bones and teeth. The most important electrolytes for human body function are the cations sodium (Na+), potassium (K+), calcium (Ca2+), magnesium (Mg2+) and hydrogen (H+), and the anions chloride (Cl–), bicarbonate (HCO3

–), hydrogen phosphate (HPO4

2–), and sulphate (SO42–).

The role of electrolytes in the body varies but they carry electrical impulses that facilitate nerve impulses and muscle contractions, contribute to body structures such as bone and teeth, regulate body fluids and control the acid-base balance of the body. Table 3-8 has been included for your interest only, and it shows some of the essential electrolytes in the body.

Table 3-8: Essential electrolytes in the human body

Electrolyte Normal conc.

Found Role Excess Deficit

sodium (Na+) 135 - 145 mmol/L.

outside body cells

Maintenance of a normal balance of body fluids. Is critical for the transmission of electrical signals to the brain and nervous system.

Hypernatraemia: thirst, oedema

Hyponatraemia: diarrhoea, decrease in body fluids

potassium (K+)

3.5 - 5.0 mmol/L.

inside of cells

regulation of the heartbeat and function of the muscles Control of body fluid and cell function

Hyperkalaemia: irritability, cardiac arrest

Hypokalaemia: lethargy, muscle weakness

chloride (Cl-) 98 - 108 mmol/L (serum)

fluid outside of cells and in blood

Maintenance of a normal balance of body fluids

Hyperchloremia: diarrhoea and overactivity of the parathyroid glands.

Hypochloremia: heavy sweating, vomiting adrenal gland and kidney disease

calcium (Ca2+)

2.45mmol/L (total plasma calcium)

Bones and blood

Used in the formation of bone and teeth, facilitates muscle contraction, involved in functioning of many enzymes, blood clotting and maintaining normal heart rhythm.

Hypercalcaemia: relaxed muscles

Hypocalcaemia: cramps, tetany

magnesium (Mg2+)

0.70 – 0.95 mmol/L (total plasma magnesium)

Bones and blood

Muscle and nerve function

Drowsiness hypertension

Adapted from Strobe, 2008. & http://www.medicinenet.com/electrolytes/page2.htm

http://www.medicinenet.com/electrolytes/page2.htm


It is important to maintain the correct balance of electrolytes in the body. During heavy exercise many of the essential electrolytes for the body, particularly sodium and potassium, are lost through sweat. These electrolytes must be replaced to keep the electrolyte concentrations of the body fluids constant. This may be done be having sports drinks which have sodium chloride or potassium chloride added to them. An excess of electrolytes can be dangerous, for example lethal injections of potassium chloride have been used for executions, causing cardiac arrest.

In Australia, the preferred unit for concentration of electrolytes are millimole per litre (mmol/L). However, American textbooks mostly use the unit milliequivalents per litre (mEq/L or mEq/litre). These units emphasise the fact that ions vary in the amount of electric charge that they carry. For example, the sodium (Na+) ion has a single positive charge, whereas the calcium (Ca2+) ion has two positive charges.

If the ions in the electrolyte have one positive or negative charge, the concentration in mEq/L is the same as that for mmol/L. For example:

Sodium (Na+) 145 mEq/L = 145 mmol/L Chloride (Cl–) 100 mEq/L = 100 mmol/L

If the ions in the electrolyte have two positive or two negative charges, the concentration in mEq/L is two times that for mmol/L. For example:

Magnesium (Mg2+) 2 mEq/L = 1 mmol/L Hydrogen phosphate (HPO4

2–) 3 mEq/L = 1.5 mmol/L

Acids and Bases We frequently encounter acids and bases in our daily lives. Oranges, lemons and vinegar are examples of acidic food. We have hydrochloric acid in our stomach to help destroy harmful microbes and digest food. Sulphuric acid is used in the batteries in our cars. As for bases, we take antacids such as milk of magnesia for heartburn, and use household ammonia as a cleaning agent.

Acids are chemical substances that produce hydrogen ions (H+ ions) when dissolved in water. Acidic solutions normally taste ‘sour’ or ‘tart’ (e.g., vinegar). Because acids form ions when they dissolve in water, they can also be considered to be a special type of electrolyte.

The names and molecular formula for some common acids are provided in Table 3-9.

Table 3-9: Names and molecular formulas for common acids.

Chemical name Molecular formula Common uses

Hydrochloric acid HCl Pool acid

Sulfuric acid H2SO4 Car batteries

Nitric acid HNO3 Fertiliser production

Acetic acid CH3COOH Vinegar

Carbonic acid H2CO3 Carbonated soft drinks


Bases are chemical substances that produce hydroxide ions (OH– ions) when dissolved in water. Again, bases can be considered to be a special type of electrolyte. Solutions of bases in water are also called ‘alkaline’ solutions or ‘basic’ solutions. These solutions have a bitter or caustic taste, and a slippery or soapy feel.

The names and molecular formula for some common bases are provided in table 3-10.

Table 3-10: Names and molecular formula for common bases.

Chemical name Molecular formula Common uses

Sodium hydroxide NaOH Caustic soda in oven cleaner

Potassium hydroxide KOH Soap manufacture

Ammonia NH3 Household cleaners

Magnesium hydroxide Mg(OH)2 Milk of magnesia for indigestion

Both acids and bases are defined as substances that dissociate to produce either hydrogen ions or hydroxide ions in solution.

Both acids and bases dissociate in water to produce ions. If dissociation (that is the reaction to produce ions) is complete, the acid or base is referred to as being a strong acid or strong base.

If dissociation is not complete, the acid or base is referred to as being a weak acid or weak base.

The following table provides examples of strong and weak acids and bases and the dissociation reactions.

Table 3-11: Strong and weak acids and bases

Classification Chemical Dissociation reaction

Strong acid hydrochloric acid + -HCl H +Cl→

Weak acid acetic acid -3 3CH COOH CH COO + H+

Strong base sodium hydroxide

+ -NaOH Na +OH→

Weak base ammonia + -3 2 4NH H O NH +OH+

Note that for the dissociation reactions the use of the single, “→”, arrow indicates complete dissociation and is used for strong acids and bases whereas the double, “ ”, arrow indicates partial dissociation and is used for weak acids and bases.


Activity 3-14

Predict whether the following solutions are acidic or basic.

Solution Acid or base? coffee

ammonia

wine

vinegar

sea water

blood

distilled water

milk

What is meant by pH? You will no doubt be familiar with the term “pH”, but what does it mean?

pH is a measure or scale of how acidic or basic a solution is. The pH scale ranges from 0 to 14 where:

pH range between 0.0 and 7.0 indicates an acidic solution

pH = 7.0 indicates a neutral solution

pH range between 7.0 and 14.0 indicates a basic (or alkaline) solution

The following diagram illustrates the full pH scale from 0 to 14, and provides the pH values for some common substances. Notice that blood has a pH of about 7.4, and urine has a pH of about 6.0. Gastric juices have a pH of about 1.6 because the stomach secretes hydrochloric acid (HCl). The HCl in gastric juice is used to activate digestive enzymes that break down dietary proteins.

Figure 3-20: The pH scale of acidity showing the pH values of some common substances (Adapted from Martini, 1998.)


Calculation of the pH of a solution is based on the concentration of hydrogen ions present in solution. There is a relationship between the concentration of hydrogen ions and hydroxide ions which allows for the calculation of pH for basic solutions.

Acid-base neutralisation As mentioned previously, the reaction of an acid with a base is referred to as a neutralisation reaction. If equal volumes of the same concentration of a strong acid and a strong base are added together, then the resulting solution becomes neutral (pH=7). This occurs since the hydrogen ions that are produced from the acid react with the hydroxide ions that are produced by the base to form pure water, according to the reaction:

+ -2H +OH H O Eqn 3-6

The products of this type of reaction are always a salt and water. Indeed, some of you may remember from your junior science days the old saying that ‘an acid plus base gives a salt plus water’.

An example of a neutralisation reaction is, when hydrochloric acid and sodium hydroxide react together. The resulting products are sodium chloride (NaCl—better known as ‘table salt’) and water (this is formed by the combination of the H+ and OH– ions, which are produced from the HCl and NaOH respectively). If both hydrochloric acid and sodium hydroxide solutions are of the same concentration, and equal volumes of each were used, both acid and base would be completely neutralised. This chemical equation for this reaction is:

2Acid Base A salt WaterHCl + NaOH NaCl + H O→ Eqn 3-7

Buffers When any amount of acid or base is added to a solution, there is a change in hydrogen ion concentration which then results in a change in pH. In some situations pH must be maintained within a narrow range. An example of this is in pool chemistry where pH must be controlled to ensure the effectiveness of chlorine, avoid skin and eye irritations and to protect the equipment and pool surface from the corrosive nature of the chemicals. Another example of application of buffers is in blood so as to ensure effective absorption of oxygen and other bodily functions.

The pH of blood needs to be maintained between 7.35 and 7.45. Outside of this range, the central nervous system and oxygen uptake of patients is affected, as shown in the following diagram.


7.8

7.7

7.6

7.5

7.4

7.3

7.2

7.1

7.0

Alkalosis

Acidosis

Normal pH 7.35 - 7.45BloodpH

Central nervous system isoverexcited; respiratoryarrest; death.

Central nervous system is depressed; coma.

Figure 3-21: The extremes of pH in our blood

A buffer contains two chemical parts—a weak acid and its conjugate base or a weak base and its conjugate acid. (A conjugate base is the substance that forms when an acid releases a hydrogen ion, and a conjugate acid is the substance that forms when a base produces hydroxide ion.) Our blood contains a very important buffer called the carbonic acid-bicarbonate buffer. This buffer contains the weak acid, carbonic acid (H2CO3) and its conjugate base, bicarbonate (HCO3

–).

Both chemical parts of a buffer play a very important role in resisting a change in pH. Together, they act like a ‘sponge’ that ‘mops up’ the small amounts of acid or base that are added to a solution.

If acid is released into the blood from tissues (lactic acid is produced by the muscles), the bicarbonate ions are able to neutralise the acid by reacting with the hydrogen ions.

i.e. - +3 2 3HCO +H H CO Eqn 3-8

If base is released into the blood, the carbonic acid is able to neutralise the base by reacting with the hydroxide ions.

i.e. - -2 3 3 2H CO +OH HCO +H O Eqn 3-9

The carbonic acid-bicarbonate buffer is therefore very effective in maintaining blood pH. Any excess carbonic acid formed is broken down according to the following equation:

2 3 2 2H CO CO +H O Eqn 3-10

The carbon dioxide (CO2) formed is then released from the body via the breathing process.


Organic compounds The element carbon has a special role in chemistry because it has the capability of forming so many different compounds. Organic chemistry is the study of all compounds that contain carbon, regardless of their origin. It is a vast area of chemistry, with well over 10 million different carbon compounds being known to chemists. Petrol, coal, medicines and perfumes are organic compounds. The fabric in your clothes may be a naturally occurring organic compound such as cotton or silk, or they may be synthetic organic compounds such as polyester and nylon.

Carbon is the basis of life on earth. The food we eat, for example, is composed of many different organic compounds that supply us with fuel for energy and with the carbon atoms needed for building and repairing the cells of our bodies.

Besides carbon, there are other elements in organic compounds, including: hydrogen, oxygen, nitrogen, sulfur, phosphorous, chlorine, bromine and iodine.

Organic compounds are grouped together into classes. Each class has a characteristic functional group. A functional group is a group of atoms within the organic compound which is responsible for most of its chemical and physical properties.

Table 3-12 shows some of the more common functional groups known to occur in organic compounds.

Table 3-12: Classification of the functional groups of organic compounds and the biomolecules in which they are found

Functional group name Structure Biomolecule(s) found in

Alkanes (all single bonds) -C-C- CnH(2n+2) Saturated fatty acids

Alkene (carbon-carbon double bond)

C C

CnH2n Unsaturated fatty acids

Alkyne (carbon-carbon triple bond)

C C CnH(2n-2) Unsaturated fatty acids

Alcohol R – OH Carbohydrates

Aldehyde R C H

O

Carbohydrates

Ketone R C R 1

O

Carbohydrates

Carboxylic acid R C OH

O

Amino acids and fatty acids

Ester

R C O R 1

O

Triglycerides

Amine R – NH2 Amino acid


For example, the characteristic functional group for alcohols consists of oxygen and a hydrogen atom, the symbols are –OH. The letter ‘R’ in a chemical formula is an abbreviation for the rest of the structure of the compound, which does not take part in a reaction. R’, indicates another carbon chain.

Some organic compounds may have two or more functional groups. Note also the terms saturated which refers to molecules with all single bonding between the carbon molecules and unsaturated where double bonds are present.

IUPAC naming of organic molecules

The name of an organic molecule is based on both the functional group that is present and the number of carbons in the longest continuous carbon chain of the molecule. Consider the following alkanes in Table3-13.

Table 3-13: Alkane names based on length of carbon chain

Number of carbon atoms in the longest carbon chain of molecule

Alkane name Molecular formula Condensed structural formula

1 Methane CH4 CH4

2 Ethane C2H6 CH3 CH3

3 Propane C3H8 CH3 CH2 CH3

4 Butane C4H10 CH3 CH2 CH2 CH3

5 Pentane C5H12 CH3 CH2 CH2 CH2CH3

6 Hexane C6H14 CH3 CH2 CH2 CH2 CH2CH3

7 Heptane C7H16 CH3 CH2 CH2 CH2 CH2CH2CH3

8 Octane C8H18 CH3 CH2 CH2 CH2 CH2 CH2 CH2CH3

Each different functional group has a light variation to the basic alkane name, and this is best explained in the following table.


Table 3-14: Naming organic compounds

Functional group name

General formula IUPAC ending

Structure based on 3 carbon chain

Name based on 3 carbon chain

Alkanes (all single bonds)

-C-C- -ane CH3 CH2 CH3 propane

Alkene (carbon-carbon double bond)

C C

-ene

CH3C CH2

H

propene

Alkyne (carbon-carbon triple bond)

C C -yne CH3C CH

propyne

Alcohol R – OH -ol CH3 CH2 CH2OH propanol

Aldehyde R C H

O

-al

CH3CH2C

O

H

propanal

Ketone R C R 1

O

-one

CH3CCH3

O

propanone

Carboxylic acid R C OH

O

-oic acid

CH3CH2C

O

OH

propanoic acid

Ester R C O R 1

O

alkyl -oate

CH3CH2C

O

OCH3

methylpropanoate

Amine R – NH2 -amine CH3 CH2 CH2 – NH2 propanamine


Activity 3-15

Complete the following table:

Condensed structural or molecular formula

Functional group

IUPAC name

CH3 CH2 CH2 CH2 CH2OH

CH3CH2CH2CH2CH2C

O

H

CH3CH2CH2C CH

CH3C

O

OH

CH3CH2CH2CH2CH2CH2CH2NH2

CH3C

O

CH2CH3

CH3CH2C CH2


Biomolecules

Biomolecules are organic molecules that are produced in living organisms. Important biomolecules include carbohydrates, lipids, amino acids and proteins.

Carbohydrates Carbohydrates are composed of the elements; carbon, hydrogen and oxygen. (The name means hydrates of carbon, or water associated with the carbon atom– i.e. CH2O) Carbohydrates are classified chemically into three major types:

1. Monosaccharides (meaning one sugar) which are the simple sugars. The important saccharides in human nutrition include glucose, galactose and fructose.

2. Disaccharides (meaning two sugars) which are composed of two monosaccharide units. The important disaccharides in human nutrition include sucrose, lactose and maltose.

3. Polysaccharides (meaning many sugars) which are composed of many monosaccharides. The important polysaccharides in human nutrition include starch, glycogen and cellulose.

Lipids The name lipid refers to any components of living cells that are not water soluble. The lipids are a diverse group of compounds in structure and function. The best known lipids are fats and oils. Less well known lipids include phospholipids, cholesterol, steroids and prostaglandins.

Fats and oils are the most prevalent form of lipids and include butterfat, lard, body fat and the vegetable oils. They are a type of ester, also known as triglycerides

The phospholipids are a class of lipids containing glycerol, two fatty acids and a phosphate group. The phospholipid molecule has both polar and non-polar parts which contribute to the effective function of the cell membrane, (this will be discussed in Unit 4). The polar part (called the polar head) is comprised of the phosphate and glycerol. The nonpolar part is comprised of the two fatty acids, which are sometimes referred to as the two nonpolar tails.

The steroids are a class of lipids that include cholesterol, bile salts, vitamin D, sex hormones, and hormones secreted from the adrenal cortex.


Amino acids and proteins Proteins are very large molecules. The building blocks of proteins are much smaller molecules called amino acids, (amino, indicating many acids). All of the 20 common amino acids have the basic structure of a central carbon atom, which is bonded to:

• an amine (–NH2) group

• a carboxylic acid (–COOH) group

• a hydrogen (–H)

• an R group (–R) which represents the rest of the molecule.

This is illustrated as follows:

Aminegroup

Carboxylic acid

Centralcarbon H

R C COOH

NH2

Figure 3-22: Representation of an amino acid All amino acids are identical except for the group of atoms called the R group. Differences in the numbers and arrangement of the atoms making up the R group make each amino acid unique in its chemical structure and properties.

Proteins are composed of many amino acid groups linked together by “peptide bonds” and can be represented as shown in Figure 3-23.

Figure 3-23: Representation of a protein

A special group of proteins are called enzymes. These proteins act as catalysts. Recall that catalysts were discussed in the section on chemical reactions and energy in reactions.


Bibliography Gilbert TR, Kirss RV & Davies G, 2004, Chemistry the science in context, WW Norton & Company Inc, New York,

Hickman, R & Caon, M 1995, Nursing science, 2nd edn, Macmillan, Melbourne

Leong, J (ed), 2008, ‘Learn about liver’, The Helix, Issue 119, April-May 2008, CSIRO Education, Dickson

Rowlands, P., 2005, ‘Plastic fantastic’, The Helix, Issue 102, June-July 2005, CSIRO Education, Dickson

Strob Martini, FH 1998, Fundamentals of anatomy and physiology, 4th edn, Prentice-Hall, Upper Saddle River, New Jersey, U.S.A.e, P,2008, Bodyworks Physics and Chemistry for Health Students,Pearson Australia

Timberlake, K 1992, Chemistry, 5th edn, Harper Collins, New York

Baker, S. & Worthley, The essentials of calcium, magnesium and phosphate metabolism: Part 1. Physiology, viewed 5th December, 2008 http://www.anzca.edu.au/jficm/resources/ccr/2002/december/CPMI.pdf

Johnson, L., Magnesium, viewed 5th December 2008, http://www.merck.com/mmhe/sec12/ch155/ch155h.html

MedicineNet, Electrolytes, viewed 5th December, 20082008http://www.medicinenet.com/electrolytes/article.htm

http://www.anzca.edu.au/jficm/resources/ccr/2002/december/CPMI.pdf

http://www.merck.com/mmhe/sec12/ch155/ch155h.html

http://www.medicinenet.com/electrolytes/article.htm

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