Preliminary Chemistry Study Notes (part 2)

WATER

TABLE OF CONTENTS

Water ...starts page 39

- Course Outcomes and Topics ... 39

OUTCOMES:

8.4.1. WATER ON EARTH ... STARTS PAGE 41 - Define the terms solute, solvent and solution ... 41

- Identify the importance of water as a solvent ... 41-42 - Compare the state, percentage and distribution in water in biosphere, lithosphere, hydrosphere and atmosphere ... 41

- Outline the significance of water and different states of water on Earth in terms of water as: a constituent of cells and its role as both a solvent and a raw material in metabolism, a habitat in which temperature extremes are less than nearby terrestrial habitats, an agent of weathering of rocks both as a liquid or solid, a natural resource for humans and other organisms ... 42

8.4.2. STRUCTURE AND PROPERTIES OF WATER ... STARTS PAGE 43 - Construct Lewis electron dot structures of water, ammonia and hydrogen sulfide to identify the distribution of electrons ... 14, 43 - Compare molecular structure of water, ammonia and hydrogen sulfide, the differences in their molecular shapes and in their melting and boiling points ... 43, 45 - Describe hydrogen bonding between molecules ... 45

- Identify the water molecule as a polar molecule ... 44

- Describe the attractive forces between polar molecules as dipole-dipole forces ... 45

- Explain the properties of water in terms of its intermolecular forces: surface tension, viscosity, boiling and melting points ... 43-46 8.4.3. WATER AS AN SOLVENT ... STARTS PAGE 47 - Explain changes, if any, to particles and account for those changes when the following types of chemicals interact with water: a soluble ionic compounds such as sodium chloride, a soluble molecular compounds such as sucrose, a soluble or partially soluble molecular element or compound such as iodine, oxygen or hydrogen chloride, a covalent network substance such as silicon dioxide ... 47-48 - Analyse the relationship between the solubility of substances in water and the polar nature of the molecule ... 47

8.4.4. DISSOLUTION AND CONCENTRATION ... STARTS PAGE 48 - Identify some combinations of solutions will produce precipitates, using solubility data ... 48

- Describe a model that traces the movement of ions when solution and precipitation occur ... 49

- Identify the dynamic nature of ion movement in a saturated dissolution ... 49

- Describe the molarity of a solution as the number of moles of solute per litre of solution ... 50

8.4.5. HEAT CAPACITY ... STARTS PAGE 51

- Explain what is meant by the specific heat capacity of a substance ... 51

- Compare the specific heat capacity of water with a range of other solvents ... 51

- Explain and use the equation for change in enthalpy ... 52

- Describe dissolutions which release heat as exothermic and give examples ... 52-53 - Describe dissolutions which release heat as endothermic and give examples... 52-53 - Explain why water’s ability to absorb heat is important to aquatic organisms and to life on Earth generally ... 53

- Explain what is meant by thermal pollution and discuss the implications for life if a body of water is affected by thermal pollution ... 53-54

TOPICS:

8.4.1. WATER ON EARTH ... STARTS PAGE 41 - Distribution of water on Earth ... 41

- Solutions... 41

- Water and natural processes ... 42

- Density ... 43

8.4.2. STRUCTURE AND PROPERTIES OF WATER ... STARTS PAGE 43 - Structure of water and other molecules ... 43

- Polar and non-polar molecules ... 44

- Intermolecular forces ... 44

- Properties of water ... 45

8.4.3. WATER AS AN SOLVENT ... STARTS PAGE 47 - Uses of water as an solvent ... 47

- Ionic substances in solution ... 47

- Molecular substances in solution ... 47

- Covalent network substances in solution ... 48

8.4.4. DISSOLUTION AND CONCENTRATION ... STARTS PAGE 48 - Precipitation reactions ... 48

- Precipitation and ion interactions ... 49

- Equilibrium ... 49

- Concentration ... 49

- Molarity ... 50

- Calculations with solutions ... 50

8.4.5. HEAT CAPACITY ... STARTS PAGE 51 - Specific heat capacity ... 51

- Enthalpy ... 52

- Heat changes when substances dissolve ... 53

- Consequences of thermal properties of water ... 53

8.4.1 WATER ON EARTH

DISTRIBUTION OF WATER ON EARTH

The presence of water on Earth in the solid, liquid and gaseous states is unique within the solar system. The fact that the Earth has retained this water is a result of the planet’s gravitational force, determined by its mass and the Earth’s temperature, determined by its distance from the Sun. Water plays a central role in the evolution of life on Earth, and remains critical to the maintenance of life on the planet.

Water has some remarkable properties that are crucial to the maintenance of all living systems. Such properties include:  Water is the only pure substance on Earth commonly found in all three states (solid, liquid and gaseous).

 Ice is less dense than water at 0°C, so as a result, ice floats allowing many aquatic life-forms to survive below the surface in cold climates.

 Water is transparent to visible light, which allows photosynthesis to occur in aqueous systems.

 Water has a high surface tension, which is important in some physiological processes and also capillary action in soils and plants.

 Water is an excellent solvent, allowing biological processes to occur in aqueous solutions, and also allows water to serve as a transportation system for nutrients and waste products in living organisms.

WATER IN THE SPHERES

 97% of the water on the planet is present in the hydrosphere. Liquid water covers nearly 75% of the Earth’s surface in the form of oceans, seas, lakes, ground water and rivers, which are collectively referred to as the hydrosphere.

 The polar ice caps cover another portion of planet Earth, dominating the regions north of the Arctic Circle and south of the Antarctic Circle.

 The atmosphere also contains varying amounts of water vapour (ranging from 0–5%).

 Water in the lithosphere can occur as moisture or permafrost within soils and rocks or as water of crystallisation within minerals. Waterofcrystallisation (or waterofhydration) refers to the water molecules that form part of the crystal structure of many ionic substances, such as hydrated copper (II) sulfate – CuSO4.5H2O

Water is the major constituent of living matter, representing approximately 70% of the biosphere. It is a component of all cells, the transport system for nutrients and waste products in living organisms, a reaction medium for many biochemical processes. Together with carbon dioxide, water is a reactant in photosynthesis.

SOLUTIONS

The oceans and seas dominate planet Earth. They are not pure substances, but rather mixtures containing large quantities of dissolved substances, called solutions. Solutions are defined as homogenous mixtures. In a solution, one substance, called the solute is dissolved in another substance, called the solvent.

AQUEOUS SOLUTIONS

Solution in which the solute is dissolved in water, are called aqueoussolutions. Common naturally occurring aqueous solutions include sea water, blood plasma, stomach acid, sap in plants and drinks like coffee and lemonade.

Aqueous solutions are indicated in chemical equations using the (aq) solution to the right of the solute.

SATURATED AND UNSATURATED SOLUTIONS

If sugar is continually added to water, there comes a point at which no more will dissolve. At this stage the solution is said to be saturated. A saturated solution is one in which no more solute will dissolve under the existing conditions of temperatures and pressure. An unsaturatedsolution is a solution that contains less than the quantity of solute needed to saturate it under existing conditions.

SOLUBILITY

Different substances can exhibit very different solubilities in a particular solvent. For example, salt (sodium chloride) is an example of a substance that is readily soluble in water, while oxygen is less soluble and sand (mainly silicon dioxide) is virtually insoluble.

WATER AND NATURAL PROCESSES

WATER AND THE BIOSPHER

 Water makes up approximately 70% of the biosphere and has a crucial role in the maintenance of life on Earth.  Water is an essential reactant in the production of glucose by the process of photosynthesis:

6CO2 (g) + 6H2O(l) + light energy → C6H12O6 (aq) + 6O2 (g)

 Glucose is the form of energy required for living organisms to function.

 Cellular respiration is essentially the reverse process, and provides the energy to sustain life. C6H12O6 (aq) + 6O2 (g) → 6CO2 (g) + 6H2O(l) + energy

 Water acts a solvent for many solutes. Water dissolves oxygen, various salts and nutrient, and is therefore important in supporting plant and animal life.

 Many waste products such as carbon dioxide, ammonia and urea are soluble in water, providing an important mechanism by which plant and animal-wastes can be removed.

 The human body is over two-thirds water, with blood, being over 80% water.  Water is also a major component of the lymph system.

 Water also is part of the moisture lining of the lungs, allowing the gases oxygen and carbon dioxide to be transferred between the air breathed in, and the respiratory system.

 Water is used to transfer heat around the body to cells performing respiration to our skin, where it can escape.

WATER AS A TEMPERATURE MODERATOR

 Water has a relatively high heat capacity, which implies it can absorb and thus release large quantities of heat with relatively small temperature changes.

 Water affects the temperature experienced in coastal regions.

WATER AND ICE AS AN AGENT OF WEATHERING AND EROSION

 Water performs a vital role in both landform formation, and soil production through the natural processes of weathering and erosion.

 Water along with dissolved acids and oxygen gas attack minerals breaking then into smaller components.

 When water freezes in cracks of rocks, its forms ice. Since ice has a larger volume than liquid water, it expands causing cracks and the breakdown of the rock in smaller fragments. This process is called icewedging.

WATER AND ITS SIGNIFICANCE

Water is essential to human survival. Access to a reliable supply of clean quality water is a necessity for modern-day society. Water has numerous uses, including drinking, washing, and in industry, especially agriculture. Water is used as a transport system, for goods and services around the globe. Water also can serve as a source of entertainment, leisure, relaxation and enjoyment.

DENSITY

 As the temperature increases, it usually expands, resulting in an increased volume and therefore a decreased density. This is due to the increasing kinetic energy of the particles in the substance.

 Most substances increase in density when they freeze, again, because a decrease of volume occurs during the freezing process.

DENSITY OF WATER

 Water is unusual in that its density that firstly it expands from the liquid to the solid state, resulting in a decreased density. The lower density of ice in relation to liquid water allows ice to float on top of water, which is crucial for the survival of life forms in polar environments and other regions that experience extremely cold temperatures.

 As water is cooled, its density increases until 4°C, after which its density decreases. This occurs because water initially contracts with cooling but at temperatures below 4°C it expands due to intermolecular forces in water causing the water causing the water molecules to arrange themselves in a particular fashion.

8.4.2 STRUCTURE AND PROPERTIES OF WATER

STRUCTURE OF WATER AND OTHER MOLECULES

Water consists of an oxygen atom covalently bonded to two hydrogen atoms. The Lewis Electron Dot Diagram does not tell us anything about the shape of molecules. It just tells us how the valence electrons are arranged and how many covalent bonds are formed.

Some electron pairs are involved in bond formation; others that are not involved are non-bonding pairs or lone pairs. Electron pairs repel each other and move as far apart as possible to achieve greater stability. The most stable way to arrange four pairs of valence electrons is in the shape of a tetrahedron. Methane, ammonia and water are molecules based on the tetrahedron.

𝜎 =m V

Calculating density:

where σ is density in kg.m–3_{, m is mass in kilograms (kg), V is volume in cubic metres (m}3₎ Note: If mass is in grams (g) and volume is in cubic centimetres cm3_{,then the density is in g.cm}3

Water has its maximum density of 1 gm.cm–3 _{at 4°C.}

Different shapes of different compounds:

 Methane (CH4) is a perfect tetrahedron

 Ammonia consists of three bonding pairs and one pair and is said to be pyramidal.  Water has two bonding electron pairs and two non-bonding electron pairs.  Hydrogen sulfide is similar to water.

O

δ–

H

δ+

H

δ+

N

δ–

H

δ+

H

δ+ δ+

H

O

C

δ–

C

δ–

C

δ–

H

δ+

H

δ+ δ+

H

H

δ+

POLAR AND NON-POLAR MOLECULES

 Elements attract electrons according to their electronegativity.

 Elements have differing electronegativity, so some elements attract electrons more than others.  Electrons spend more time with the stronger elements than weaker elements.

 This results in one end of a molecule being more negative than the other end.

A pair of equal and opposite charges separated in space is called a dipole. Covalent bonds in which electrons are unequally shared are called polarcovalentbonds. Polar molecules have a net dipole. Non polar molecules have their electrons shared equally. For example: H2, Cl2, and N2 are non-polar molecules since, the electrons are shared equally but HCl and H2O are HCl is very electronegative. Electrons are shared unevenly.

COMMON POLAR BONDS

 Common polar bonds are the bonds between hydrogen and oxygen, hydrogen and nitrogen, carbon and oxygen, carbon and chlorine.

 Common non-polar bonds are the bonds between nitrogen and chorine, carbon and sulfur, phosphorus and hydrogen, carbon and hydrogen (only slightly).

INTERMOLECULAR FORCES

 The forces of attraction between molecules are called intermolecularforces.

 The physical properties of covalent substances are dependent on the strength of these forces.

 The chemical properties of these molecules are dependent mainly on the strength of the chemical covalent bonds which hold the atoms of the molecule together.

DISPERSION FORCES

 Dispersionforces are present in all covalent molecular substances.

 Dispersion forces are weak attractive forces which result from uneven electron distribution within atoms or between neighbouring atoms and molecules

 Dispersion forces can be visualised as the attraction between temporary dipole and induced temporary dipoles.

DIPOLE-DIPOLE FORCES

 Dipole-dipoleforces are present in some covalent molecular substances, when the component atoms have different electronegativities.

 Since polar molecules have positive and negative ends they are able to line up so that the positive end of one molecule attracts the negative end of another molecule.

 Electrostatic forces of attraction hold the molecules strongly together.  Dipole-dipole attractions between molecules raises the melting point

and boiling point of these substances for example: HCl and H2S

Diatomic molecules that have a polar covalent bond are called polarmolecules. Polyatomic molecules are not necessarily polar, because two or more dipoles can cancel each other out. The shape determines whether dipoles cancel or add up to give a net dipole.

In the case of water and ammonia these dipoles add up, but in the case of carbon dioxide (because it is linear) and methane (because it is a tetrahedron), these dipoles cancel.

Order of Electronegativity:

HYDROGEN BONDING

 Hydrogenbonding is a strong form of the dipole-dipole force and is present in some covalent molecular substances containing hydrogen with fluorine, oxygen or nitrogen atoms.

 Hydrogen bonds are particularly strong types of polar interactions, as result when a hydrogen atom bound to one of three atoms (fluorine, oxygen or nitrogen) can become attached to another one of these atoms in another molecule. Water (H2O) and ammonia (NH3) display hydrogen bonding, as well as molecules such as H2O2, H2SO4, CH3OH and other alcohols.

 Hydrogen bonds are much stronger than ordinary dipole-dipole forces. Their relative strengths arise from the fact that the hydrogen nucleus (a bare proton only) is extremely small and that the fluorine, oxygen and nitrogen strongly attract.

PROPERTIES OF WATER

MELTING POINT AND BOILING POINT OF WATER

 With consideration to metals, ionic substances and covalent lattices, water has a relatively low melting point and boiling point. This is expected since water is a covalent molecular compound.

 However, compared to similar molecules, the melting point and boiling point of water appears to be far higher than expected.

Molecular Formula Molecular Mass Melting Point (°C) Boiling Point (°C)

CH4 16 -182 -160

NH3 17 -78 -33

H2O 18 0 100

HF 20 -83 19

Molecular Formula Molecular Mass Melting Point (°C) Boiling Point (°C)

H2O 18 0 100

H2S 34 -83 -62

H2Se 81 -66 -40

H2Te 130 -49 -2

These strange properties of water can be explained by the hydrogen bonding between molecules of water.  Water molecules are held together by dispersion forces, dipole-dipole forces and also strong hydrogen bonds.  Ammonia molecules are also held together by the same forces as water molecules, but the hydrogen bonds in

ammonia are not as strong as the hydrogen bonds in water, because nitrogen is not as electronegative as oxygen.  Intermolecular forces for hydrogen sulfide are weaker because of the lack of hydrogen bonding, resulting in a low

melting point and boiling point.

 As a result more energy is required to separate water molecules from each other resulting in a higher melting and boiling point.

DENSITY

 In water, the molecules are held together by hydrogen bonds and fit relatively close together.

 However, in ice, the hydrogen bonds force the molecules further apart as they form a three dimensional crystal lattice.  As a result, ice has a larger volume than liquid water, resulting in ice having a lower density than liquid water.

SURFACE TENSION

 The surfacetension of a liquid is a measure of the elastic forces in the surface layer.

 A molecule that is inside a liquid, experiences intermolecular forces from the other molecules around it.  A molecule in the surface layer experiences intermolecular forces only from molecules beside it and below it.  Surface tension is due to the unbalanced downward forces from other molecules acting on molecules at the surface.  The stronger the intermolecular forces, the greater will be the surface tension of the liquid.

 Water which has strong hydrogen bonding has strong intermolecular forces and so a high surface tension, since the water molecules at the surface experience stronger forces than in comparable liquids.

Liquid Surface Tension (mNm–1₎

mercury 487 water 72.6 chloroform 27.1 acetone 23.7 methanol 22.6 ethanol 22.3 hexane 18.4

VISCOSITY

Viscosity is a measure of a liquid’s resistance to flow or being poured. Viscosity can be measured by comparing the rate of flow of liquids. The higher the viscosity, the slower the rate of flow is.

There are two factors that affect the ease with which molecules can move over one another: One factor is the size and complexity of the molecules themselves. For example, Motor oil contains very long floppy molecules that get tangled up and do not flow easily. The other factor is intermolecular forces. The stronger the forces of attraction between pairs of molecules, the more resistance there is to flow.

Syrup and honey are solutions of sucrose, glucose and fructose in water. All these four compounds contain O-H groups and so there is extensive hydrogen bonding between the molecules and water. This gives rise to the high viscosity of syrup, honey and glycerol.

Water seems to have a low viscosity, however when we compare it with many other pure liquids such as ethanol, acetone chloroform and hexane its viscosity is comparatively high because there is extensive hydrogen bonding involving all the atoms of the molecule.

 Viscosity depends on the forces between the molecules of the liquid.  The forces between water molecules are strong hydrogen bonds.

 Water with its strong hydrogen bonding results in a much higher resistance to flow than its small molecular size might suggest. As a result water has a high viscosity when compared to similar pure liquids.

Liquid Viscosity (mPas)

glycerol 1490 mercury 1.53 water 0.891 chloroform 0.58 hexane 0.33 pentane 0.224

8.4.3 WATER AS AN SOLVENT

USES OF WATERS AS A SOLVENT

Water is the most widely used solvent by both humans and natures. It has many usages including:  chemical reactions

 dissolving nutrients

DETERMING DISSOLUTION

Solvents are either polar or non-polar solvents. Knowing this can help chemists determine whether dissolution will occur if two or more of these substances are mixed together. Water is known as the universal solvent since it can dissolve a wide range of substances. Solutions formed with water as the solvent, are called aqueoussolutions.

IONIC SUBSTANCES IN SOLUTION

 Most ionic substances dissolve in water, because polar water molecules have a strong attraction to charged ions.  The ionic bonds in the ionic substance break in the crystal lattice break apart. The ions are move freely and

independently. The hydrogen bonds between the water also break

 The positive ends of the water molecules attach to the negative ions, and the negative ends of the water molecules attach to the positive ions. This happens because the attraction between water and the ions are stronger than the attractive forces between the positive ions and negative ions.

 Since free ions exist in the solution, the solution is therefore able to conduct electricity and is called an electrolyte.

 Some substances crystallise out with molecules of water still attached, called waterofcrystallisation.

MOLECULAR SUBSTANCES IN SOLUTION

SOLUBLE MOLECULAR SUBSTANCES

 The crystals of molecular substances that are soluble will break up and disperse throughout the solid and break up into individual molecules.

 Molecules will only dissolve only if the water molecules form a stronger attachment to the solute molecules than the attachment of solute to solute.

 Polar substances will dissolve in water (which is also a polar substance), but non-polar substances will not dissolve.  Some molecular substances will dissolve in water by reacting with the water molecules. This often results in the

formation of ions (this process is called an ionisationreaction), resulting in the formation of an acid or base.

PARTIALLY SOLUBLE MOLECULAR SUBSTANCES

 Some molecular substances are usually elements that exist as molecules, such as iodine or oxygen. These substances will only partially dissolve in water. These substances are non-polar substances, but are partially soluble because of the weak dispersion interactions with the solvent.

 Solute-water interactions are so weak that solubility is low, and as a result dissolution does not to completion.

INSOLUBLE MOLECULAR SUBSTANCES

 The very large molecules of some hydrocarbons, such as polymers will mean the presence of strong hydrogen bonds.  Water molecules are unable to break these large molecules and so these substances will not dissolve.

Polar solvents dissolve polar solvents. Non-polar solvents dissolve non-polar solvents.

COVALENT NETWORK SUBSTANCES IN SOLUTION

 Strong covalent bonds exist throughout the entire solid.

 As a result, water molecules are unable to break these strong covalent bonds and therefore these substances will not dissolve.

8.4.4 DISSOLUTION AND CONCENTRATION

PRECIPITATION REACTIONS

A precipitation reaction is when a solid is produced from a reaction of two soluble compounds, resulting in the

production of an insoluble solid which is suspended in the solution. Precipitation occurs because not all ionic compounds are soluble and that results in an insoluble compound being precipitated, if the separate ions are added together in solution. It is assumed that if the precipitation reaction goes to completion that either one or both of the ions involved will be completely precipitated.

DETERMINING SOLUBILITY

WRITING IONIC EQUATIONS

Often chemical equations involve ions, rather than neutral molecules, which require ionicequations to be properly represented. Precipitation reactions are best represented using ionic equations.

Some disadvantages of using a neutral species equation to represent precipitation reactions include:

 The neutral species equation implies that before the reaction take place the ions are in some way stuck together, i.e. bonded together. However, what is actually happening is the ions exist separately in water.

 The equation implies that two of the ions (the spectator ions) have changed in some way during the reaction. However, these ions are called spectator ions because they do not take part in the chemical reaction. They still exist as ions dissolved in the solution.

The best way to represent precipitation reactions is to: 1. Ions are written separately when they are in solution

2. Spectator ions which, although they may be present, do not take part in the reaction are not included (these are called spectator ions)

3. Numbers of atoms and charge must be balanced.

With a precipitation reaction it is important to always use state symbols, as this indicates which compound is the precipitate, and which compound still exists dissolved in the solution as ions.

Soluble compounds:

 All nitrates  All acetates

 All sulfates (except those of calcium, barium, lead, mercury and silver).  All chlorides, bromides and iodides (except those of lead, mercury and silver)  All ammonium compounds

 All sodium, potassium (and the rest of Group 1) compounds. Insoluble compounds:

 All phosphates, sulfites, carbonates (except those of Group 1 elements and ammonium)

 All oxides and hydroxides (except those of Group 1 elements, calcium, magnesium, barium and ammonium)  All sulfides (except those of Group 1 elements, Group 2 elements and ammonium)

PRECIPITATION AND ION INTERACTIONS

DISSOLUTION WITH INSOLUBLE COMPOUNDS

A saturated solution is a solution where no more solute will dissolve into the solvent at existing conditions. If more solute tries to be dissolved in results in the presence of undissolved solid. Saturation, the limit of dissolving the solute,

is reached earlier with an insoluble compound than a soluble compound.

When an insoluble ionic compound is added to a solvent such as water, a small amount of the crystal lattice will break into free ions. When the solution reaches saturation, no more of the compound will dissolve and some of the ionic compound remains undissolved in the solution.

ION MOVEMENT IN PRECIPITATION

Insoluble compounds will precipitate from solutions. Precipitation will occur when the quantity of insoluble ions exceeds the solubility of the compound formed from these ions.

Precipitation is essentially the opposite process to dissolution (dissolving). In a saturated solution, the ions in the solid are continually leaving the crystal to go into the solution, while at the same time, and at the same rate, different ions in the solution precipitate out, resulting in neither reaction going to completion.

EQUILIBRIUM

In saturated solutions, a situation called equilibrium occurs when precipitation and dissolution occurs at the same time and at the same rate resulting in there being no apparent activity. During this situation, the total concentration of ions and the amount of precipitate will remain constant.

In order to set up such a situation, the containing vessel must be sealed to prevent evaporation of the solvent and the temperature needs to be kept constant.

CONCENTRATION

Concentration is defined as the amount of solute in a specific amount of solvent. This is expressed in many ways, because:  Concentration is used by people other than chemists, who may be unfamiliar with the concept of molarity.

 In commerce and industry, the measure of concentration emphasises the mass or volume of the solute, rather than the moles.

 In areas such as environmental studies or drug testing, the concentration of the solute is very small, so molarity is impractical, and thus concentration is used instead. The concentration used here is usually parts per million (ppm) Some common ways of expressing concentration are:

 Mass of solute per volume of solvent or solution.  Volume of solute per volume of solvent or solution.  Mass per volume as a percentage – %(w/v)  Volume per volume as a percentage – %(v/v)  Mass per mass as a percentage – %(w/w)

 Parts per million (ppm), mass in milligrams per kilogram of solution

Equilibrium is a dynamic situation in a closed system, where there is continual interchange between reactants and products on an atomic level, with no noticeable change in observation or physical properties.

Equilibrium Equations:

MOLARITY

The concentration of a solution in moles per litre indicates the number of moles of solute dissolved in a litre of solution. This form of concentration is the most commonly used in chemistry, as it allows from the measurement of a definite volume of solution of known concentration, the number of moles of solvent present in the solution.

The concentration of a solution in moles per litre is referred to as molarconcentration or molarity. It is sometimes given the symbol M. (A 0.50 mol L–1 _{solution is the same as a 0.50 M solution)}

PREPARING AND DILUTING SOLUTIONS

PREPARING SOLUTIONS OF KNOWN CONCENTRATION

 A weighed mass of substance is dissolved in water and the solution made up to a definite volume in a volumetric flask.  The beaker is washed out to ensure that all the solute is transferred to the volumetric flask. If this is not done

properly, then the concentration will be different. The volumetric flask needs to be filled to the marked level. If this is not done the concentration will be different and also harder to calculate.

 The expected solution can be calculated using the formula for calculating concentration from moles.

DILUTING SOLUTIONS

 First, a solution of known concentration is prepared.

 A sample (determined by the needed dilution) is then taken and transferred to another volumetric flask, which also needs to be filled up to the marked level

CALCULATIONS WITH SOLUTIONS

c =n

V

Calculating molarity:

where C is the concentration in mol L–1 _{(moles per litre), n is number of moles and V is volume in litres}

Calculating dilution: C1V1 = C2V2

where C1 is the original concentration and C2 is the final concentration in mol L–1 _{(moles per litre), V1 is the original} volume and V2 is the final volume in litres.

8.4.5 HEAT CAPACITY

SPECIFIC HEAT CAPACITY

 Water plays an important role that water plays, particular the oceans, in moderating temperatures. This is because of the relatively large quantities of heat absorbed or released when the temperature of water changes.

 This demonstrates the fact that water, unlike other substances such as metals or concretes, does not increase in temperature as fast when heated.

 For example a bucket of water when exposed to same amount of radiant energy from the Sun as the surrounding environment does not increase in temperature to the same extent as the surroundings.

 The explanation for this that water has a specific heat capacity than the surroundings.

SPECIFIC HEAT CAPACITY OF VARIOUS SUBSTANCES

Substance _{(kJ kg}Specific Heat Capacity _{–1 K–1 or J g–1 K–1)} Substance _{(kJ kg}Specific Heat Capacity _{–1 K–1 or J g–1 K–1)}

water 4.18 aluminium 0.90

pentane 1.66 chloroform 0.55

ethanol 1.41 carbon tetrachloride 0.54

toluene (methylbenzene) 1.13 glass 0.50

phenol (hydroxybenzene) 1.11 iron 0.45

benzene 1.05 copper 0.39

nitrogen gas 1.04 silver 0.23

oxygen gas 0.92 mercury 0.14

CALCULATING HEAT ABSORBED OR RELEASED

The specific heat capacity may be used to determine the energy absorbed or released, when a temperature of a known mass or substance changes.

CALORIMETRY

 In order to measure heat changes during a chemical reaction, we use a calorimeter.

 If two objects are brought into contact, heat will flow from the hot object to the cold object until the temperature of the two objects are equal.

 Two objects can be at same temperature but contain different amounts of heat

The specificheatcapacity (C), also called specific heat, is the amount of energy required to change the temperature of 1 gram of substance of a substance by 1 Kelvin (or Celsius).

Q = mC∆T

Q = mC∆T

Calculating the amount of heat absorbed or released:

where Q is the amount of heat released or absorbed in kilojoules, m is mass in kilograms, C is the specific heat capacity, and ΔT is the temperature change in Kelvin (or Celsius)

where Q is the amount of heat released or absorbed in joules, m is mass in grams, C is the specific heat capacity, and ΔT is the temperature change in Kelvin (or Celsius)

 Express the final answer in a positive number. Use the words released or absorbed to specify whether the heat  Make sure that the final answer in in kilojoules not joules, i.e. for the second formula, remember to divide by 1000  If using Q for further calculations, such as enthalpy or molar heat, make sure to remember sign

ENTHALPY

The release of energy in chemical reactions occurs when the reactants have a higher chemical energy than the products. The chemical energy of a substance is a type of potential energy stored within the substance. This stored chemical potential energy is called the heat content or enthalpy of the substance is given the symbol H.

CALCULATING CHANGE IN ENTHALPY

EXOTHERMIC REACTIONS

 If the enthalpy in the system decreases during a chemical reaction, a corresponding amount of energy (Q) must be released to the surroundings, i.e. the enthalpy of the products is less than the enthalpy of the reactants.

 The enthalpy difference between the reactants and the products is equal to the energy released to the surroundings.  A reaction is which heat energy (Q) is released to the surroundings is called an exothermicreaction. Examples of

exothermic reactions include synthesis reactions.

 Since exothermic reactions release energy into the surroundings, the result of this type or reaction is that the surroundings heats up.

ENDOTHERMIC REACTIONS

 If the enthalpy in the system increases during a chemical reaction, a corresponding amount of energy (Q) must have been absorbed from the surroundings i.e. the enthalpy of the products is greater than the enthalpy of the reactants.  The enthalpy difference between the reactants and the products is equal to the energy absorbed from the

surroundings.

 A reaction is which heat energy (Q) is absorbed from the surroundings is called an endothermicreaction.  Since endothermic reactions absorb energy from the surroundings, the result of this type of reactions is that the

surroundings are cooled down.

 Decomposition reactions are endothermic reactions, as often the energy input required for the reaction is absorbed from the surroundings.

 A change in the enthalpy of a substance is given the symbol ΔH.

 It is the difference between the total enthalpy of products and the enthalpy of the reactants, i.e. the change in the enthalpy of a system. It is equal to the following expression: ∆H = ΣH products − ΣH reactants

Calculating the change in enthalpy: ∆H = −mC∆T

where ΔH is change in enthalpy, m is mass in g, C is the specific heat capacity of the substance being cooled or heated, and ΔT is change in temperature in Kelvin (or degrees Celsius)

Important things to remember:

 Always convert the final answer to kilojoules, and where specified: kilojoules per mole, which are the units usually used for ΔH

 Always pay careful attention to the sign. Exothermic reactions should be a negative answer, while endothermic reactions should be a positive answer.

 If kilojoules per mole are required, use relevant formula to convert mass to moles and then divide through by the number of moles to the get the change in enthalpy per mole.

HEAT CHANGES WHEN SUBSTANCES DISSOLVE

When ionic substances dissolve in water, there is a noticeable change in temperature. This means the reaction is endothermic (absorbing heat) or exothermic (releasing heat). For example:

i. When sodium hydroxide NaOH dissolves in water, the solution heats up. The dissolution process releases heat which warms up the solution. The dissolution of NaOH is said to be exothermic.

ii. When potassium nitrate KNO3 dissolves in water, the solution cools. It requires an input of energy which is taken from normal thermal energy of the water and the solid substance. The dissolution of KNO3 is endothermic. Energy is needed to break the ionic bonds in the crystal lattice of the solute, and energy is also needed to break the intermolecular forces (i.e. the hydrogen bonding) between water molecules. But energy is released when the separated ions form bonds with water molecules. Determining these factors will help determine the whether the dissolution is exothermic or endothermic.

MOLAR HEAT OF SOLUTION

If ΔHsoln is positive, then the reaction is endothermic, but if ΔHsoln is negative, then the reaction is exothermic.

CONSEQUENCES OF THERMAL PROPERTIES OF WATER

Living organisms can survive and reproduce only if their temperatures are maintained within fairly narrow ranges. Water within cells provides the necessary temperature regulation because of the following properties that water has:

 High heat capacity (large amount of heat absorbed produce small temperature rises)

 High thermal conductivity relative to other liquids (quickly removes heat from the hot location to the cooler one)  Large proportion of most living things.

The high heat capacity of water means the aquatic environment is maintained at a stable temperature, and therefore allows aquatic organisms to thrive.

It also has a moderating effect on global temperatures, and is a large factor in influencing global temperatures. It also smooths day to night, summer to winter fluctuations in temperature, and produces a more hospitable environment then more terrestrial environments.

WATER POLLUTION

Water pollution is the contamination of water bodies, for example, lakes, rivers, oceans and groundwater. Water pollution directs affects the organisms which live in these bodies of water, and has harmful effects on natural ecosystems, such as ocean ecosystems and river ecosystems.

CAUSES OF WATER POLLUTION

Water pollution is caused by chemical contaminants and pathogens. Water pollution is sometimes caused by contaminants are released from a single point such as pipe, or discharges from sewerage treatment. Water pollution is sometimes caused by accumulation of contaminants or contaminants that do not have a distinct source, and may be discharged through a pipe as stormwater or waste.

Pathogens such as salmonella and the norovirus (which causes the common cold), may result from poor treated sewerage disposed of in the sea. These pathogens could lead to water-borne disasters affecting the human and animal populations dependent on the marine or river ecosystems.

The molar heat of solution ΔHsoln of a substance is the heat absorbed when one mole of the substance dissolves in a large excess of water

Chemicals including detergents, insecticides, herbicides, petroleum and other fuels, and organic industrial solvents (VOCs) are organic compounds that act as pollutants when discharged into a water body, and are present in concentrated

amounts. Inorganic chemicals, such as ammonia, fertilizers, sediment (run-off from construction or logging), chemical waste and acids, are substances that also cause water pollution. Some chemical waste such as rubbish in the form of food waste, plastics and paper and ships cause visible damage to the water ecosystem, and create debris.

THERMAL POLLUTION

One particular type of water pollution is thermalpollution. This is the discharge into a river or lake quantities of hot water that are large enough to increase significantly the temperature of the water body. A 2-5°C increase can be significant. Some of the effects of thermal pollution include:

 Less dissolved oxygen, which results in stress for the organism

 Increase metabolism of the organism, due to increased temperature increases demand for oxygen and so aggravates the decreased dissolved oxygen problem

 Fish eggs do not develop or hatch if temperature is too high. Sudden temperature changes can kill fish eggs even if change is within survival range

 False temperatures cues can set off migration and spawning at the wrong time.  Lethal temperature can be exceeded.

Thermal pollution can result from when the water ecosystem is used for cooling in industry or electrical generation. This is a concern for electricity stations, especially those located on shallow lakes along NSW’s Central Coast. It is rarely a

problem when open ocean is used. The problem can be combatted with cooling towers so cooling water is completely recycled, or alternatively by using cooling ponds.

EFFECTS OF WATER POLLUTION

The effects of these pollutants, in addition to the effects of thermal pollution, include:  Some pollutants cause discolouration, temperature changes or changes in acidity.  Some pollutants cause cloudiness, which block the gills of certain fish species.

 Some pollutants cause oxygen depletion (anoxia) which may affect the populations of certain species. Some pollutants cause this indirectly through the increase of concentration of certain chemical nutrients which results in an increase in the productivity of the ecosystem.

 Some pollutants are toxic, killing organisms. Other pollutants, which are pathogens, can cause diseases which could reduce the population numbers of certain species.

 Some pollutants cause reductions in water quality or changes to salinity of the water.  Some pollutants cause changes in electrical conductivity.

ENERGY

TABLE OF CONTENTS

Water ...starts page 55 - Course Outcomes and Topics ... 55

OUTCOMES:

8.5.1. SOURCES OF ENERGY ... STARTS PAGE 57 - Outline the role of photosynthesis in transforming light energy to chemical energy and recall the raw materials for

this process ... 57 - Outline the role of the production of high energy carbohydrates from carbon dioxide as the important step in the

stabilisation of the sun’s energy in a form that can be used by animals as well as plants ... 57 - Identify the photosynthetic origins of chemical energy in coal, petroleum and natural gas ... 57 8.5.2. CARBON AND ITS PROPERTIES ... STARTS PAGE 58 - Identify the position of carbon in the Periodic Table and describe its electron configuration ... 58 - Describe the structure of diamond and graphite allotropes and account for their physical properties in terms of

bonding ... 58-59 - Identify that carbon can form single, double or triple covalent bonds with other carbon atoms ... 58 - Explain the relationship between carbon’s combining power and ability to form a variety of bonds and the existence

of a large number of carbon compounds. ... 58 8.5.3. HYDROCARBONS AND THEIR PROPERTIES ... STARTS PAGE 59 - Describe the use of fractional distillation to separate the components of petroleum and identify the uses of each

fraction obtained ... 59 - Identify and use the IUPAC nomenclature for describing straight-chained alkanes and alkenes from one carbon to

eight carbons ... 60-61 - Compare and contrast the properties of alkanes and alkenes carbon 1 to carbon 8 and use the term homologous

series to describe a series with the same functional group ... 60-62 - Explain the relationship between the melting point, boiling point and volatility of the above hydrocarbons, and

their non-polar nature and intermolecular forces (dispersion forces). ... 62 - Assess the safety issues associated with the storage of alkanes carbon 1 to carbon 8 in view of their weak

intermolecular forces (dispersion forces). ... 62 8.5.4. COMBUSTION ... STARTS PAGE 64 - Describe the indicators of chemical reactions ... [18, 52], 65 - Identify combustion as an exothermic chemical reaction ... [52], 64 - Outline the changes in molecules during chemical reaction combustion in terms of bond-breaking and

bond-making ... 65 - Explain that energy is required to break bonds and energy is released when bonds are formed ... [19], 65 - Describe the energy needed to begin a chemical reaction as activation energy ... 66 - Describe the energy profile for both endothermic and exothermic reactions ... [52], 65-66 - Explain the relationship between ignition temperature and activation energy ... 66 - Identify the sources of pollution which accompany the combustion of organic compounds and explain how these

can be avoided ... 64 - Describe chemical equations using full balanced equations to summarise examples of complete and incomplete

8.5.5. REACTION KINETICS... STARTS PAGE 67 - Describe combustion in terms of slow, spontaneous and explosive reactions and explain the conditions under which

these occur ... [64], 67

- Explain the importance of collisions between reacting particles as criterion for determining reaction rates ... 68

- Explain the relationship between temperature and the kinetic energy of the particles ... 69

- Describe the role of catalysts in chemical reactions, using a named industrial catalyst as an example ... 68

- Explain the role of catalysts in changing the activation energy and hence the rate of chemical reaction ... 69

TOPICS:

8.5.1. SOURCES OF ENERGY ... STARTS PAGE 57 - Photosynthesis ... 57

- Fossil Fuels ... 57

- Water and natural processes ... 42

- Density ... 43

8.5.2. CARBON AND ITS PROPERTIES ... STARTS PAGE 58 - Carbon and bonding ... 58

- Allotropes of carbon ... 58

8.5.3. HYDROCARBONS AND THEIR PROPERTIES ... STARTS PAGE 59 - Fractional distillation of crude oil ... 59

- Classifying hydrocarbons ... 60

- Molecular substances in solution ... 47

- Covalent network substances in solution ... 48

8.5.4. COMBUSTION ... STARTS PAGE 64 - Reactions with oxygen ... 64

- Combustion reactions ... 64

- Chemical reactions ... 65

- Heat of combustions ... 65

- Activation energy ... 66

8.5.5. REACTION KINETICS... STARTS PAGE 67 - Determine the rate of a reaction ... 67

- Combustion and reaction rates ... 67

- Factors affecting the rate of reaction ... 68

- Collision theory ... 68

8.5.1 SOURCES OF ENERGY

PHOTOSYNTHESIS

People in industrialised societies use large amounts of energy and most of it comes from fossil fuels. This energy comes directly or indirectly from the sun via photosynthesis.

Photosynthesis is the process in which plants use solar energy to convert carbon dioxide from the air and water from the ground into carbohydrates such as glucose, sucrose, starch and cellulose, which are compounds containing carbon, oxygen and hydrogen:

Carbon dioxide + water + light→ glucose + oxygen 6CO2 (g) + 6H2O (l) + light → C6H12O6 (aq) + 6O2 (g)  Photosynthesis is an endothermic reaction.

 It is a multistep reaction brought about by the green pigment chlorophyll  Solar energy is converted into chemical energy and is stored.

 Carbohydrates in plants are the energy source for animals, including humans.

 They are high energy compounds as they release large amounts as they release large amounts of energy during respiration or combustion.

Cellularrespiration is essentially the reverse process to photosynthesis. It converts the stored chemical energy into a form which the organism is able to then use:

Glucose + oxygen → carbon dioxide + water + energy C6H12O6 (aq) + O2 (g) → 6CO2 (g) + 6H2O (l)

FOSSIL FUELS

Normally when plants and animals die, insects, fungi, worms and bacteria called decomposers, convert them back into carbon dioxide, water and nutrients, releasing the stored energy into the environment.

In some locations and under certain conditions, over geological time scales, the decay processes were interrupted

resulting in some plant material only being partially decomposed. This partially-composed material remains stored in the form of rich compounds that are called fossil fuels. All of these compounds are carbon-based since living matter is composed of carbon.

Through the processes of mining, and burning fossil fuels, the energy content of these materials can be used as a power source for transportation and for generating electricity for homes, schools and businesses.

Some fossil fuels include:  Coal

 Crude oil  Natural gas  Oil shales  Tar sands

Australia has an abundance of fossil fuel resources, in particular its immense supplies of coal and natural gas, but has limited reserves of crude oil. As a result Australia uses lots of fossil fuels to produce electricity, and exports some of these resources overseas in particular to big consumers such as China or India.

The amount of energy released during respiration is the same as what was absorbed during photosynthesis, which is 2830 kJ per mole of glucose used.

8.5.2 CARBON AND ITS PROPERTIES

CARBON AND BONDING

Carbon is classified as a non-metal. It is the first member of group IV of the periodic table and is located in period 2 between boron and nitrogen. The carbon atom has 6 electrons and therefore has an electron configuration is 2, 4. This means it can either lose 4 electrons or gain 4 electrons to gain a stable electron configuration. It usually however tends to covalent bond with other non-metals such as hydrogen and oxygen.

Since it has 4 valence (free) electrons, each carbon atom can be bonded with 4 hydrogen atoms, or 2 oxygen atoms. Carbon atoms can therefore bind to form many compounds including hydrocarbons and carbohydrates. For this reason it a useful base for the compounds of life. The large number of compounds is due to its ability to form bonds with many elements, which can be single, double or triple bonds.

Carbon can from single bonds such as those in alkanes, double bonds such as those in alkenes, or triple bonds such as those is present in alkynes. Carbon also bonds with itself to produce different allotropes.

ALLOTROPES OF CARBON

Allotropes are forms of one element in the same physical state which have distinctly different physical properties i.e. colour, density, hardness and electrical conductivity). Diamond and graphite are some of the allotropes of carbon. Other elements which display allotropy are arsenic, phosphorus, selenium, sulfur, tin and oxygen. Allotropes have different properties because the atoms are joined or packed together in different ways to form molecules or crystals.

AMORPHOUS CARBON

 Amorphous carbon (soot) is a low pressure form of carbon that usually exists in the form of as imperfect tetrahedron structures. Each carbon is surrounded by 4 other carbons bonded to it with single bonds.

DIAMOND

 The carbon atoms in diamond are covalently bonded to four other carbon atoms to form a 3D covalent lattice, with the shape around each carbon atom being

tetrahedral. The 6 membered rings are buckled not flat.

 All valence electrons are tied up in strong covalent bonds, so there are no mobile electrons. Thus it doesn’t conduct electricity.

 It has an orderly arrangement of atoms throughout the whole crystal gives its transparency and brilliance. For this reason it is used in jewellery.

 Diamond is a hard substance because of the strong forces between molecules and it used therefore for drills and saws.

GRAPHITE

 Each carbon atom in graphite is bonded only to three other carbon atoms, forming a planar structure of flat six-membered rings.

 This leaves each carbon atom with 1 free valence electrons which forms a delocalised electron cloud located between the layers (called aromaticity), similar to that present in metals.

 These electrons are able to move within sheets. As a result, graphite is able to conduct electricity. However, it can cause a temporary dipole in one sheet which by electrical induction will cause a dipole of the opposite charge to occur in a neighbouring sheet. For these reasons, it is used for the electrodes in dry cell batteries and superconductors.  In graphite, weak intermolecular forces between layers are present, resulting in these layers being able to readily

sheer off or to slide over one another. For these reasons, it is used for pencils, and also as a dry lubricant.

 Graphite lubrication abilities are hypothesised to be due to the presence of a fluid layer (i.e. air) between the layers (as graphite is a poor lubricant in a vacuum).

OTHER ALLOTROPES OF CARBON

 Electrical discharge between graphite electrodes in low pressure helium causes graphite to evaporate and then condense as soot. If soluble in benzene or toluene, it forms a yellow-brown substance with 60 carbon atoms that has a soccer ball structure.

 Buckminster fullerenes (bucky-balls) are 5-6 membered rings that combine to form a spherical cage. The most common has 60 carbon atoms. Other fullerenes with 70, 74 and 80 carbon atoms have also been produced. In fact, the number of carbon atoms can range from 32 to 84 carbon atoms.

 Bucky-balls have some delocalised electrons but not like graphite.

 Cage-like fullerenes may have uses as superconductors or as lubricants because of the weak intermolecular forces between their ball-shaped molecules. However since they are expensive to produce, other materials are used instead.  Nanotubes (or bucky-tubes) are other group of fullerenes consist of molecules that have tube-like shape, that are

several nanometres long.

 Nanotubes have high tensile strength, and can act as either conductors or insulators

8.5.3 HYDROCARBONS AND THEIR PROPERTIES

FRACTIONAL DISTILLATION OF CRUDE OIL

Crude oil is a complex mixture of hydrocarbons formed by geological action on decayed aquatic plant and animal matter over long periods (i.e. millions of years).

The first step in oil refining is fractional distillation. For separating components of crude oil it is carried out in large steel towers up to 40 metres high. In this process, the components of oil are separated according to their boiling points. Since boiling point increases as molecular weight increases, the separation is roughly in order of increasing atomic weights (or increasing number of carbon atoms per molecule). The crude oil is then vaporised by heating, they fed into the bottom of the fractionating column which contains a series of trays.

The temperature falls as the vapour rise up through the column. The least volatile components, i.e. those with the highest boiling points and hence the largest molecular weights condense near the bottom of the columns while the most volatile components do not condense until they reach the top of the column. Liquids are drawn off from the column at various heights and these are the various fractions which are collected.

Lubricating oils are obtained from the least volatile fraction by distilling under a vacuum because they will boil off at a lower temperature that would be needed at atmospheric pressure). Greases are separated from the remaining non-volatile material by solvent extraction and the final residue is asphalt or tar.

Fraction Name Boiling Point Range (°C) Carbons Uses

natural gas <20 1-4 household gas, making methanol and hydrogen in plants petroleum ether 20-100 5-7 very inflammable solvent

benzines 70-90 6-7 safer dry cleaning solvent than carbon tetrachloride

ligroin 80-120 6-8 solvent

gasoline 40-205 5-10 motor cars (the main refinery product) kerosene 175-325 12-18 aviation and tractor fuel

diesel oil >275 13-18 diesel engine fuel for trucks and trains lubricating oil (refinery liquid) 16-20 lubricants

vaseline greases (refinery solid) 18-22 pharmaceuticals

paraffin waxes (hard solid) 20-30 candles, cartons, surf-board wax bitumen (hard solid) 30-40 roof flashing, road asphalt

CLASSIFYING HYDROCARBONS

The most important fossil fuels i.e. natural gas, petroleum and coal, are all mixtures of hydrocarbons. Hydrocarbons are compounds that contain the elements carbon and hydrogen (not to be confused with carbohydrates which contain the elements oxygen, carbon and hydrogen).

There are three main groups of hydrocarbons:  aliphatic

 alicyclic  aromatic

IUPAC (International Union of Practical and Applied Chemistry) has developed a systematic system to name hydrocarbons and other organic compounds. The purpose of adopted this naming system (nomenclature) is to establish an

international standard of naming organic compounds. Each organic compound is given a name which effectively describes its structure.

When naming organic compounds, attention must be paid to:  The number of carbon atoms in the hydrocarbon chain  The presence of any functional groups in the compounds.

ALKANES



Alkanes are described as saturated hydrocarbons since each carbon is bonded to four other atoms, meaning that no other atoms can be incorporated into its structure. In all alkane molecules, each carbon atom forms four single bonds.  When naming alkanes, the suffix –ane is used.

No of Carbons Name No of Carbons Name No of Carbons Name

1 methane 8 octane 20 icosane

2 ethane 9 nonane 21 henicosane

3 propane 10 decane 22 docosane

4 butane 11 undecane 30 triacontane

5 pentane 12 dodecane 31 hentriacontane

6 hexane 13 tridecane 40 tetracontane

7 heptane 14 tetradecane 50 pentacontane



Methane (CH4), Ethane (C2H6), Propane (C3H8) and Butane (C4H10) are gases at room temperature. The boiling point increases with chain length, so the next twelve alkanes are liquids and the rest are solid at room temperature.



Alkanes are found in natural gas and crude oil. Natural gas is mostly methane, with small amounts of ethane, propane and butane. Crude oil is a much more complicated mixture of hydrocarbons, and can contain alkanes with up to 100 carbon atoms in their molecules.



Alkanes are unreactive, apart from combustion. For example acids and alkalis have no effect on them. However they do burn well in a good supply of oxygen, forming carbon dioxide and water vapour. The reactions give out plenty of heat, so alkanes are often used as fuels. When propane burns, the reaction is:

C3H8 (g) + 5O2 (g) → 3CO2 (g) + 4H2O (g) + heat



Although alkanes do undergo substitution reactions with some very reactive halogens such as fluorine and chlorine.



Both propane and butane are used as camping gas, and in gas lighters. Calor gas is mainly butane, and natural gas is

used for cooking and heating in homes.

Homologousseries are groups of related chemical that obey a general formula and share the same functional groups, i.e. particular aspects of their structure, such as double bond or triple.

In alkanes, the general formula is CnH(2n+2) where n is an integer.

ALKENES

 Alkenes are unsaturated hydrocarbons that contain at least one carbon double bond. The suffix –ene is used to indicate the double bond present. The double bond in alkenes is the functional group.

 For alkenes containing four of more carbon atoms, structural isomerisation occurs. Isomers have the same molecular formula but are arranged differently. When naming alkenes, the longest chain must contain the double bond, and is numbered so that double bond has the lowest number possible.

No of Carbons Name No of Carbons Name No of Carbons Name

1 n/a 3 propene 5 pentene

2 ethene 4 butene 6 hexene

 Alkanes also burn in oxygen like alkanes, and have very similar chemical properties. e.g. C2 (g) + 3O2 (g) → 2CO2 (g) + 2H2O(g) + heat

 However, unlike alkanes which are relatively unreactive compounds, alkenes are not, reacting with hydrogen and other compounds, e.g. C2H4 (g) + H2 (g) → C2H6 (g)

 Alkenes are more reactive than alkanes. The double bond within carbon atoms can break to form single bonds, so can combine easily with other elements such as hydrogen or oxygen.

 Ethene is unsaturated since its molecules can add on more atoms, while ethane is saturated since its molecules cannot fit in more atoms because there are no double bonds to break, and each carbon atom already has four single bonds.  Alkene molecules can combine with each other, due to their double bonds in a process called polymerisation. During

polymerisation, many small molecules, called monomers, join together to from very large molecules, called polymers.  Polythene is an example of such a molecule which is formed by this process. It is formed, when ethene is heated, under

high pressure. More than 1,000 ethene molecules can combine through this process to make a single molecule of polythene. Polythene is a solid. It is unreactive, as there are no double bonds present. It can be rolled into thin sheets and moulded into different shapes, and because that is easy to mould, it is called a plastic.

There are two ways of testing a hydrocarbon to determine whether it is an alkane or alkene.

1. The addition of bromine water. Bromine water is an orange solution of bromine in water. It turns colourless in the presence of an alkene, because the bromine adds on to the alkene, to form a colourless compound:

C2H4 (g) + Br2 (aq) → C2H4Br2 ethane + bromine water → 1,2-dibromoethane

2. The addition of potassium manganate. This is purple, but turns colourless when an alkene is present.

ALKYNES

 Alkynes are unsaturated hydrocarbons that contain at least one carbon triple bond. The suffix –yne is used to indicate the triple present.

 When naming alkynes, the longest chain must contain the triple bond, and is numbered so that triple bond has the lowest number possible.

BRANCHED GROUPS

 When an alkane (or alkene etc.) has a branched group which replaces one of the hydrogen atoms, then the longest chain in the molecule still gives the basic name of the compound.

 The chain is numbered to give any double bonds or triple bonds as before, but where possible number so that the branch has the lowest number possible (note: the presence of a double or triple bond takes precedence).

 Name the group joined to the chain, and state the number of the carbon atom to which it is joined.

 Alkyl groups are similar to alkanes instead they have one less hydrogen, for example: methane (CH4) becomes methyl (CH3). The suffix for alkyl groups are –yl. For a chain with multiple branches, both branches are named together. Prefixes should also be added (such as di-, tri- and tetra-).

 If there are two or more different branches, the groups are written alphabetically.

 When numbering, in the case of a conflict, the one that comes first alphabetically takes precedence provided there are no other constituents.

Branch Name Structure Branch Name Structure

bromo Br– fluoro F–

PROPERTIES OF ALKANES AND ALKENES

 Alkanes and alkenes are both non-polar substances. The bonds between the molecules are weak intermolecular forces. As a result these substances have low melting and boiling point, and many of them exist as gases or liquids at room temperature. Generally the longer the carbon chain, however, the higher the melting and boiling points, which is which methane is gas, while octane is a liquid. This is because there are more bonds to break in a longer carbon chain, which means more energy is needed.

 Alkenes tend to have lower melting and boiling points than alkanes because alkenes are smaller, and therefore fewer bonds that need to be broken, and therefore less energy required to change its state from solid to liquid, or liquid to gas.

 The density naturally increases as the carbon chain increases (i.e. the molecular weight increases). Alkanes and alkenes, when present as a liquid are less than water. They also are insoluble in water because they are non-polar substance and water prefers being bonded to its own atoms (as it is a polar substance). This means when placed in water, alkanes and alkenes will float on-top of the water, forming an immiscible layer.

 Alkanes and alkenes have a high volatility, which means they go from a liquid to gaseous state much easier than low volatile liquids. Volatility also means a relatively high partial pressure of their gas above the liquid’s surface.

FLASH POINTS

The flashpoint is the lowest temperature where if a small flame is applied to the vapour above the fuel will cause it to ignite in a closed container with air. Chemicals that have a low flash point, such as hydrocarbons have low boiling points. In fact, all hydrocarbons up to octane have a flash point less the room temperature meaning that potential to ignite at room temperature (25°C).

USES OF HYDROCARBONS

The major uses of hydrocarbons are as fuels such as:

 natural gas for domestic uses, such as cooking and heating, and also in industry  LPG for caravans and barbeques (and even some cars)

 petrol and diesel for cars, trucks, vans and trucks  kerosene for jet aircraft

Higher molecular weight hydrocarbons are used as lubricating oils and greases. Vaseline is a highly purified form of these high molecular weight hydrocarbons, while paraffin waxes are a mixture of high molecular weight alkanes. Low molecular weight hydrocarbons are used as propellants in aerosol sprays. Many hydrocarbons are solvents as well.

SAFETY ISSUES WITH HYDROCARBONS

Hydrocarbons are extremely flammable (especially the low molecular weight), and are toxic in high concentrations. They also have a high volatility, meaning they can quickly form a flammable and explosive mixture with air.

As a result the following safety precautions are undertaking when using hydrocarbons:  A well-maintained cylinders and fittings for gaseous hydrocarbons are used  These cylinders are stored away from sources of electrical sparks

 Odours are added for early detection of leaks  Sturdy containers are used for liquids

 The quantities of these substances used are minimised, especially in everyday usage.  These substances are kept away from exposed flames and sparks.

 Warning signs are placed around the storage of these liquids as well as on the containers which store these substances.

 These substances should never be handled in confined spaces, instead well-ventilated rooms or a fume-hood should be used.