CONCEPT DIAGRAM (�“Mind Map�”) OF TOPICSome students find that memorizing the OUTLINE of a topic helps them learn and remember the concepts andimportant facts. As you proceed through the topic, come back to this page regularly to see how each bit fits thewhole. At the end of the notes you will find a blank version of this �“Mind Map�” to practise on.
DynamicEquilibrium
in aSaturated Solution
Water in the �“Spheres�” of the EarthIn the Atmosphere, water is present as water vapour, andas tiny liquid droplets in the clouds.
In percentage terms, water makes up between 1% and 5%of the air, varying with time, place and weather.
In the Lithosphere, water makes up about 10% of the solidEarth. Although solid rock and minerals seem perfectly dry,there is often water incorporated into the crystal latticestructure of many minerals. When volcanoes erupt, a hugeamount of steam is released as the rocks are melted.
The Hydrosphere is, of course, nearly all water. Oceanscontain about 3% dissolved salt, but the ice caps, rivers andlakes are virtually 100% water.
In Living Things, water makes up about 75% of every life-form, but jellyfish or watermelons are more like 95%
The Many Roles of Water on EarthWater is essential in all living things because it is
�• a solvent for all the chemicals in a living cell, and the medium in which all the chemical reactions occur.
�• a reactant or product in many biological reactions, such asphotosynthesis and cellular respiration.
�• a transport medium for carrying substances, such as whenfood, oxygen, etc. are carried in the blood.
�• a shock-absorber and support structure. Many plants and simple animals (e.g. worms) rely on water pressure in theirtissues to hold their body in shape. Our brain and otherbody parts are cushioned by water-based body fluids.
�• a habitat (place to live) for many species. Water habitatshave very stable temperatures because of water�’s abilityto absorb heat with little temperature change.
Water is a major factor in global climate, weather and theshaping of landforms.
�• The �“water-cycle�” produces all rain, hail and snowfall.�• Water is the main agent of erosion, carving out the valleysand wearing down the mountains, creating the landscapes.
�• Water can absorb, transport and release vast amounts ofheat energy. The ocean currents largely control globalclimates by re-distributing heat world-wide.
For humans and their civilization, water is a major resource
�• for drinking, cooking, washing and recreation.�• for crop irrigation and farming.�• in industry as a solvent, cleaning agent and cooling agent.�• for transport by boat and ship.�• for generating hydro-electricity.
A SOLUTION is a mixture, usually of a solid (the�“SOLUTE�”) and a liquid (the �“SOLVENT�”).
The solute and solvent particles are intimately associatedso that the mixture cannot be separated by filtration, andthe solute will never settle to the bottom. We say thesolute is �“dissolved�” in the solvent.
Basic Properties of WaterYou may have done some simple practical work toinvestigate some of the basic properties of water.
Water�’s Density �“Anomaly�”For almost every pure substances the solid is more densethan the liquid.
Water is the opposite... liquid water has a higher densitythan ice. How do we explain this?
In solid water (ice) the molecules form a �“molecularlattice�”. Each molecule is held rigidly in place.
When ice melts to form liquid water, the molecules haveenough energy to move around freely. However, they arestill very close together, and in fact they �“wriggle�” in evencloser to each other than when rigidly arranged in the solidlattice. Now there is the same mass of particles crammedinto less space... higher density.
Since solid ice has a lower density, it floats in liquid water.
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Density of Liquid & Solid WaterDensity is the ratio between the mass of a substance andthe space (volume) it occupies. All pure substances havea fixed and characteristic density.
Density changes with temperature.Water achieves its highest density at 4oC.
This value is 1.00 g/mL
Density = Mass Volume
D = m V
A simple method isto weigh an empty,dry measuringcylinder, then fillwith water. Read thevolume of wateraccurately then re-weigh to get themass of water.
Ideally, you wouldrepeat thesemeasurements withdifferent volumesof water.
For ice, you need to weigh it quickly before it melts. Ifthe ice cubes really are cubes or rectangular prisms, youmight measure length, width and height, then calculatethe volume.
Typical Results
Liquid water: Mass = 245 g Volume = 250mL
D = m = 245 = 0.98 g/mLV 250
Ice: Mass = 33 g Volume = 36mL
D = m = 33 = 0.92 g/cm3
V 36
Note:When measuring volume, we normally measure liquids inmillilitres (mL) and solids in cubic centimetres (cm3). Forpractical purposes these are equal volumes.
Melting & Boiling PointsPure water melts at 0oC, and boils at 100oC, undernormal �“1 atmosphere�” of pressure.
(As you may know, the celsius temperature scale is basedon the m.p. & b.p. of water.)
Under different pressures, or if impure, the m.p. and b.p.will change. For example, it can be difficult to get a good,hot cup of tea on a high mountain, because at the lowerair pressure the water boils at a much lower temperature.
You may have done experiments to find out the effect ofimpurities on the boiling point. A common experiment isto boil water with, and without, an additive such as saltand measure the boiling temperature. It will usually befound that the boiling temperature rises by severaldegrees with solute dissolved in it.
Water the Weirdo!We are so familiar with the everyday properties of waterthat we do not realize how unusual and strange water is,until we make a careful comparison with other, similarcompounds.
Some of these properties will be studied in this topic, buthere�’s a preview:
The Strange Properties of Water �• Abnormally high m.p. and b.p.�• Abnormally high viscosity and surface tension�• Abnormally high Heat Capacity�• Unusual Density anomaly (already described)
... when compared to similar sized molecules.
Why? It�’s all a matter of bonding...
Bonding in Molecular Compounds of Hydrogen
To understand water, we need to compare it to other,similar sized, covalent molecules containing hydrogen:
Methane CH4 Structural formula
Lewis Formula
The Lewis Formula, and the structural formula, wouldsuggest that the molecule is a flat, box-shape. However, thepairs of electrons in each covalent bond always try to get asfar away from each other as possible, and in 3-dimensionsthis results in a tetrahedron shape (a regular, triangularpyramid with 4 points as far apart as possible).
Each point of a tetrahedron is as faraway from the other 3as it can get.
In the methane molecule, eachcovalent bond (and therefore eachhydrogen atom) is as far away fromthe other 3 as it can get.
Ammonia NH3 Structural formula
Lewis Formula
In this molecule, the 4 pairs ofelectrons surrounding the nitrogen atom are also at the points of a tetrahedron.
However, one pair is notinvolved in a covalentbond... it is an �“unbondedpair�”, but still occupies apoint of the tetrahedron.
The result is that the ammonia molecule is a triangular pyramid shape.
Water H2O Structural formula
Lewis Formula
In the water molecule there are twounbonded pairs occupying 2 of the points of the tetrahedron.
Therefore, the water molecule is bent.The diagrams suggest a 90o right anglebetween the hydrogen atoms, but in 3-D it is more like 105o.
Hydrogen sulfide H2Shas exactly the same bonding geometry as water. Thecentral sulfur atom is larger than oxygen, but otherwisethe molecules are very similar.
The 4 compounds CH4, NH3, H2O and H2S are ofcomparable size and bonding. Now we compare theirmelting and boiling points, and how these are related totheir relative �“molecular weights�”.
Compound Molecular m.p. b.p.Weight (oC) (oC)
CH4 16 -183 -162
NH3 17 -78 -33
H2O 18 0 100
H2S 34 -86 -60
Usually, the m.p.�’s & b.p.�’s of comparable substances showa steady increase as the atomic or molecular weightincreases.
This graph shows that both water and ammonia haveunusually high melting and boiling points. Waterespecially has values way above those of comparablemolecules.
Why? What�’s going on?
In Topic 1 you learned how the properties of m.p. & b.p.are controlled by the bonding within substances.
Covalent molecules are held together internally by strongcovalent bonds (�“intra-molecular bonds�”). Thesehowever, are not the bonds that must be overcome to meltor boil the substance.
It�’s the forces between the molecules (�“inter-molecularbonds�”) that must be overcome to melt or boil a molecularsubstance.In water, it seems these forces are unusually strong!
Polar Covalent BondingTo understand water better, you must learn more aboutcovalent and ionic chemical bonding.
Up to this point, you have seen these types of bonding asquite different things. Now you must realize that they arereally different degrees of the same thing.
An analogy might help... Imagine sharing some lollies withanother person. If both of you are very fair about it, andneither dominates or intimidates the other, the sharing willbe equal:
An ionic bond can be thought of as the lolly-sharingbetween a hungry bully and a wimp who hates lolliesanyway:
Now you must learn that there is also a situation (or awhole heap of situations) in between these extremes, wherethe lollies will be shared, but perhaps not evenly.
Sharing,but notevenly.
In chemical bonding, this kind of sharing is called a �“PolarCovalent Bond�” and occurs when electrons are sharedbetween 2 atoms with quite different values forElectronegativity. (This was introduced in Topic 2... revise)
A �“Pure�” Covalent Bondoccurs when electrons areshared evenly.
In a �“Polar Covalent Bond�”the sharing is not even.The electrons areattracted more to oneatom than the other. This causes the bond (and perhaps the entire molecule) tobecome electrically �“polarized�”. The electric charge is not evenlydistributed. One end has a greater concentration of electrons andhas a slight negative charge (!!""), while the other end becomesslightly positive (!!++). The Greek letter delta (!!) is used to denotea �“small amount�” of something, in this case electric charge. Themolecule is called a �“dipole�”, meaning it has 2 poles.
Polar Bonds Create Inter-Molecular ForcesThe charges on each end of a molecular dipole are only afraction of the size of the charges on an ion, but they docause electrical forces to occur between nearby molecules.
These forces are called
Dipole-DipoleForces
It is these forces which are the �“inter-molecular forces�”that hold the molecules together in the solid state. Theseare the forces which must be overcome with thermalenergy in order for the solid to melt. These are the forceswhich determine the m.p. and b.p. of a molecularsubstance.
The strength of the dipole-dipole force varies according tothe degree of polarity of the covalent bond (how evenly orunevenly the electrons are being shared) and also variesaccording to the shape of the molecule. In somesubstances the forces are very weak, in others quite strong.
The strongest dipole-dipole forces are about 1/3 as strongas a full-scale ionic bond. These occur whenever hydrogenatoms are bonded to Oxygen, Nitrogen or Fluorine, andare called...
Hydrogen BondsOxygen, Nitrogen and Fluorine are all small, stronglyelectronegative atoms. Hydrogen is even smaller, and oncethe electrons are �“sucked away�” from it in the polar bond,the hydrogen atom is really a �“naked�” proton.
The result is an especially strong set of partial charges, apowerful dipole, and strong inter-molecular force, whichattracts nearby molecules to each other. These especiallystrong dipole-dipole attractions are called
�“Hydrogen Bonds�”.
Hydrogen Bonding in WaterIn the water molecule the covalent bonds are very polar, sothe atoms develop especially large partial charges. Eachmolecule is a dipole, and strong inter-molecular �“HydrogenBonds�” attracts each molecule to its neighbours.
It is this network of hydrogen bonds that holds themolecules in a rigid lattice in the solid state.
The Hydrogen Bonding is the reason that ice has sucha high melting point, compared to other molecules.(Ammonia also has relatively high m.p. & b.p... samereason!)
Once melted to a liquid, the molecules can move around,but �“cling�” to each other because of the hydrogen bonds.The molecules even �“wriggle�” closer to each other and thedensity increases.
To boil water to a gas, the molecules must be able to totallybreak free from the hydrogen bonds. This requiresconsiderable energy, so water has an unusually high boilingpoint, compared to other molecules.
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!! ++!! ++
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!! ++!! ""
!! ++
!! ""
!! ++
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!! ++!! ""
!! ++!! ""
!! ++!! ""
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O, N or F atomH atom
Polar Covalent Bond Hydrogen Bond
Intra-molecularCovalent Bondswithin molecules
Inter-molecularHydrogen Bonds
betweenmolecules
It is the HYDROGEN BONDINGbetween water molecules
which explains all of water�’sweird and unusual properties
More of Water�’s Unusual PropertiesAs well as the �“Density Anomaly�” and the very high m.p.and b.p., water has other properties which, compared toother similar size molecules, are quite extraordinary.
You may have done simple practical work to demonstratethese:
Surface Tensionis a phenomenon where a liquid acts as if it has a �“skin�” atthe surface. In most liquids the effect is small, but water hasa relatively strong surface tension.
Technically, the metal is NOT floating.
The explanation is, as usual, hydrogen bonding. Watermolecules have a network of forces attracting them to eachother. At the surface, this network of force resistspenetration and can support objects which will sink ifpushed through the �“skin�”.
Surface tension is also the reason that water forms droplets.
The surface tension network of forces tries to pull thedroplet into a spherical shape. The dew-drops in the photoare hanging on a spider web.
Viscosityis another phenomenon you may have experimented with.
Viscosity is a measure of how �“sticky�” or �“thick�” a liquidis. Technically it is measured as the resistance of a liquid toflowing through a thin tube, but it can be thought of ashow easy or difficult it is for things to move through theliquid.
You may have dropped marbles intovarious liquids and compared therates at which they fell, as a way toobserve viscosity differences.
Liquids like oil are very viscous, soyou may get the idea that water has alow viscosity. Yes it has, comparedto oil, but that�’s not really a faircomparison.
In fact, when the viscosity of wateris compared to liquids with similarsize molecules, water�’s viscosity isvery high.
Why? It�’s that hydrogen bonding again...
The hydrogen bonds between watermolecules cause them to �“cling�” toeach other, and make it much moredifficult for a moving object tomove through the liquid.
The high viscosity of water has had a major impact on theevolution of any aquatic animals who need to movequickly to catch food or escape predators.
Fast moving aquatic animals are always streamlined inshape and equipped with powerful tails or flippers forpropulsion.
Worksheet 1Part A Fill in the blanksIn the atmosphere, water is present as a).................................and as tiny liquid droplets in the b)................................... Inthe lithosphere, water makes up about c)..............% of solidrock, embedded as part of the d).............................................structure of many minerals. The e)...........................................is almost 100% water, most of it liquid in the f)....................,but some as solid ice in the polar g).......................................and glaciers. In living things, water makes up about h)......%of every living cell.
Water is essential to all living things because:�• it is the i).................................... for all the chemicals of life.�• it is involved in many important j).........................................�• it k)................................... substances around the body.�• it l)............................ and ............................. body organs.�• it is a m)................................. for plants & animals to live in.Water environments have very n)........................................temperatures, because of water�’s ability to o)............................................................. without much temperature change.
Water is important in controlling p)............................... andweather, and is a major agent of q)................................
For human society, water is a major resource for�• drinking, r)............................, washing and s)..........................�• for crop t).................................. in farming.
9
Emmaus Catholic College SL#802440
�• in industry, as a u).........................., and cleaning andcooling agent.
�• for generating v)........................................................
At normal atmospheric pressure, pure water melts atw).........oC and boils at x)..........oC. Its density isy)...............g/mL, but (very unusually) the density of ice isz).......................... (more/less) than liquid water. This isbecause in the liquid, the molecules actually getaa)............................ to each other than when rigidly lined upin the molecular ab)............................. of ice.
Part B Practice Problems Density D = m/V1. Complete the table of valuesSample Mass Volume Density
2. a) Which 2 samples in the table might be the same substance? Explain your answer.
b) Which substance in the table would float in water?Explain.
Worksheet 2Fill in the blank spaces. Check answers at the back.
Water has a number of unusual a)...................................including abnormally high m.p & b.p. and the �“densityanomaly�”. These are all due to the b).....................................between the molecules.
When an atom has 4 pairs of electrons around it (as is thecase in most covalent molecules) each pair tries to stayc)..................................................................... as possible. Theresult is that each pair lies at one of the points of ad)................................................. This is why methane is atetrahedral shaped molecule. In ammonia, the centrale)................................ atom is bonded to 3 f).............................atoms but also has an g)....................................... electron pairoccupying one point of the tetrahedron. Therefore, themolecule is a h).......................................... shape. In water, theoxygen atom has 2 pairs of i).....................................................occupying 2 points of the tetrahedron. This results in themolecule being j)........................ (shape)
In a �“pure�” covalent bond, the electrons are sharedk)................................... An ionic bond occurs when the�“sharing�” is totally uneven so that ions form. In betweenthese extremes there are �“l).............................. covalentbonds�” in which the sharing is m).................................. Theresult is that n)............................ electric charges (denoted bythe greek letter �“!�”) are produced on the molecule becauseof the uneven distribution of electrons.
The molecule is said to be a o).........................................because it has 2 electric poles.
The small charges on the o).............................. are not aslarge as the charge on an ion, but do create forces ofp)........................................... between each molecule and itsneighbours. These q)....................... - ............................. forcestend to hold molecules together. These are the r)..............-molecular forces which must be overcome for a substanceto change s)...........................
When hydrogen is bonded to atoms of t)...............................,..................................... or .................................... the forces areespecially strong. These are called �“u)......................................Bonds�”.
Water is such a molecule. The molecules are stronglyattracted to each other by the u)............................. bonds.This means that the m.p. & b.p. are abnormally v)..................compared to similar sized molecules.
Another result of the hydrogen bonding is that water has avery strong w)................................................. which acts like a�“skin�” and can support small objects which willx)................................... if placed under the surface.
Water also has a relatively high y).................................... dueto the way the molecules cling to each other. Because ofthis, many aquatic animals are z)........................................ toallow easier movement through the water.
Water as a SolventPerhaps the main reason that water is so important to livingthings, and in the study of Chemistry is that it is a greatsolvent.
This doesn�’t mean that everything will dissolve in water...far from it. You may have done experimental work to try tofind any general rules about which substances will, or willnot, dissolve in water. Generally, it all depends on the typeof bonding within the substance.
Ionic Compoundsare (generally) soluble in water, and all because watermolecules are polar.
Ionic compounds are composed of a strong ionic crystallattice. It requires a high temperature to melt this lattice, butwater molecules can dissolve the crystal by surroundingeach ion and detaching it from the lattice.
Notice how the the (+ve) ions are surrounded by watermolecules which are presenting the (!!"") end of theirdipole to the ion. The (-ve) ions are surrounded bymolecules presenting the (!!++) end of the dipole.
With each ion surrounded by dozens of water molecules,the attraction between the ions is �“blanketed�” and theindividual ions can no longer get close enough to eachother for their charges to bond them together.
An ionic compound in solution is made up of free moving,separate, hydrated ions.
Covalent Molecular Substancesmay, or may not dissolve, in water depending on their ownpolar nature, and on how large the molecules are.
If the solute molecules are themselves polar, they willgenerally dissolve, because the water molecules willsurround each molecule, attracted by dipole-dipole forces.
In the case of ethanol (CH3CH2OH) (alcohol) the watermolecules form hydrogen bonds with the ethanolmolecules which contain the highly polar -OH chemicalgroup.
There are many covalent molecules like this, with -NH or -OH groups on the molecule, including all the alcohols andthe �“sugars�” such as sucrose (table sugar).
Small, non-polar covalent molecules such as iodine (I2),oxygen (O2) and nitrogen (N2) will dissolve in water, butonly in small amounts... we say they are �“sparingly soluble�”.These molecules do not have any dipole charges to attracta water molecule and become �“hydrated�”, but they are sosmall and have such small �“dispersion forces�” holdingthem to each other, that they can simply spread out, insmall numbers, among the water molecules.
Larger non-polar molecules will NOT dissolve in water.They are too large to simply disperse among the watermolecules, and there are no dipoles for the water moleculesto associate with or form hydrogen bonds.
These substances include petrol, oils and waxes, and areoften described as �“hydrophobic�” (= water hating/fearing)because they will not mix with water.
In the pure state, the compound hydrogen chloride (HCl)is composed of small polar molecules:
Despite the dipole-dipole attractions, the m.p. & b.p. arequite low, so pure HCl is a gas at room temperature.
You would expect that these molecules would dissolve inwater, but they do much more than just dissolve... theyinteract so strongly with water that the molecules ionizeand become separate H+ and Cl- ions.
Hydrogen chloride dissolved in water is, of course,hydrochloric acid. This is more than just dissolving inwater because the molecule has ionized... what was a polarcovalent bond has become ionic, due the the influence ofthe polar water molecule.
This equation describes the dissolving of HCl gas toform hydrochloric acid.
Water as a Solvent (continued)
Covalent Network Substanceslike the elements Silicon and Carbon, and compounds likesilicon dioxide SiO2 (the mineral silica), are crystal latticesof atoms bonded together covalently.
Since the bonds are non-polar, or only slightly polar, watermolecules are not attracted, and the substance will NOTdissolve.
Compounds with Very Large MoleculesLiving cells produce many very large molecules, eachcontaining perhaps tens of thousands of atoms. Some, likecellulose (in plant cell walls) contain many polar groups,and water molecules will be attracted and form hydrogenbonds. However, the cellulose molecules are often linkedtogether by their own hydrogen bonding, and covalent�“cross-linking�”, and it is impossible for the huge moleculesto be taken into solution.
Cellulose is therefore insoluble, but is described as being�“hydrophilic�” (= water loving) because water will cling to it,wet it and soak into it very well.
Some protein molecules will dissolve if they have a folded,�“globular�” shape that allows water molecules to surroundthem. This is the case with enzyme proteins, which aredissolved in the water inside a cell, or in the blood.
Other proteins, like keratin (in hair and skin) are in longchains that cross-link to others. They will not dissolve, butare hydrophilic.
Plastics, such as polyethylene, are composed of hugemolecules too. Most are non-polar, and may be cross-linkedwith each other. They tend to be insoluble in water and aregenerally hydrophobic.
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!!""!!++
!!""!!++
!!""!!++
!!""!!++HCl molecules
Cl- ++
!!++
!!++!!""!!""
!!""
!!""
!!++
!!++
HCl((gg))
molecule
Separate, hydrated ionsCl-((aaqq)) and H++
((aaqq))
HCl(g) H+(aq) + Cl-(aq)
(aq) means �“aqueous�”.This is Latin for �“in water�”.In an equation it meansdissolved and hydrated by watermolecules.
Molecules in the gasstate.
�“Like dissolves like�”...water is polar, so it dissolves:�• ionic compounds�• polar molecules (unless too large)�• very small non-polar molecules (sparingly)
Ionic SolutionsWhen an ionic compound dissolves in water, the crystallattice disintegrates and the (+ve) and (-ve) ions becomeseparately hydrated to form the solution.
The positive (+ve) ions are collectively called �“cations�”.Negative (-ve) ions are known as �“anions�”.
You need to be able to write an equation to describe thedissolving of any ionic compound.More examples:
Notice that the equation must balance in terms of the ratioof the ions. In this case there are 2 nitrate ions for eachmagnesium ion. Notice also that the total of (+ve) chargesis the same as the total of (-ve) charges.
Try the Worksheet at the end of this section
Dilute, Concentrated, SaturatedIf you dissolved a pinch of salt in a bucket of water this isa �“dilute�” solution, meaning that it contains very littlesolute compared to the amount of solvent.
If you dissolved a heaped spoonful of salt in a glass ofwater the solution is �“concentrated�”... it has quite a lot ofsolute compared to the amount of solvent.
There is a limit to how much solute can be dissolved in agiven amount of solvent. When this limit is reached, andthe solution contains as much solute as it can hold, it is saidto be �“saturated�”.
Different compounds have different solubilities, and thiscan change with temperature, but as an example, at 25oC asalt-water solution is saturated when about 36g of salt havedissolved in each 100mL of water. We say the solubility ofsalt is 36 g/100mL, or simply 36 % m/v.
(�“% m/v�” means �“percentage mass to volume�” and refersto the measurement of grams (mass) in 100mL (volume).
This is not the only way we can measure the concentrationof a solution... the Mole is Back!! (soon)
Dynamic Equilibrium in a Saturated SolutionIf you keep adding and stirring salt into water until thesolution is saturated, you reach a �“dynamic equilibrium�”between the ions still in an undissolved, solid, crystal lattice,and those in the solution as separate, hydrated ions.
For simplicity in this diagram, the water molecules havebeen left out.
Since dissolving and precipitating occur at the same rate,the concentration of the solution does not change, and theamount of undissolved solid remains the same. At themacroscopic level, it seems that nothing is happening, butdown at the atomic level things are moving... ionsconstantly dissolving into solution and precipitating backout of it again. This is known as a �“DynamicEquilibrium�”
This double-arrow symbol indicates that the reaction isoccurring in both directions, at the same rate, in dynamicequilibrium.
Precipitation ReactionsNot all ionic compounds are as soluble as salt. Some reachsaturation at such a low concentration that you canconsider them as being �“insoluble�”.
Silver nitrate (AgNO3) is soluble:
Sodium chloride is soluble:
If you mix these 2 solutions together, you are really mixingwater containing 4 separate ions... Na+, Cl-, Ag+ & NO3
-.
However, silver chloride (AgCl) has an extremely lowsolubility, so the mixture of ions may contain Ag+ ions andCl- ions at concentrations way above the saturationconcentration of AgCl. The ions will immediately form anionic crystal lattice and solid AgCl will precipitate from thesolution, until the correct dynamic equilibrium of solid andsolution is re-established.
This is an ionic equation describing exactly what happened.On the left is the mixture of ions that were broughttogether in the 2 solutions. The Ag+ and Cl- ions havecombined to form solid AgCl, while the other 2 ions havestayed in solution, unchanged... they are �“spectators�”.
We can leave out the spectators to see the essential changethat occurred:
This is a net ionic equation.
Notice that it is simply the reverse of the equation for thedissolving of silver chloride.
Ionic equations can be tricky to balance. If insoluble PbCl2 isformed by precipitation of ions, the net ionic equation is:
Notice that 2 Cl- ions are needed. If these were delivered in asodium chloride solution, then to balance everything, 2 Na+
ions must be present in the full ionic equation.
You may have done experimental work as suggested by thisphoto, to discover any patterns regarding which ions areoften involved in precipitation reactions, and which mostlystay in solution.
The results of such experiments are often summarized by alist of �“Solubility Rules�”. In keeping with the K.I.S.S.Principle, here is a simplified version:
If you learn these �“rules�” you can predict what will happenwhen 2 ionic solutions are mixed:
Example 1Mix solutions of barium hydroxide & potassium iodide.
Prediction: No reaction. There is no combination of any ofthese ions which will form an insoluble precipitate.
Example 2Mix solutions of potassium carbonate with copper(II)sulfate.
Prediction: A precipitate of copper(II) carbonate will form.
Measuring Concentrations With the MoleEarlier, the idea of measuring the concentration of asolution was introduced. One way to do this is to measurethe mass of solute in each 100mL of solution (%m/v).However, although this is fairly common, it is not thestandard way to express or measure concentrations.
The Mole is Back!For reasons that will become obvious later, the standardmethod for measuring concentrations of solutions is inmoles per litre (molL-1).
Try the Worksheet, at the end of Section
Why are There Different Concentration Measurements?
Simple: it�’s a matter of convenience, for the particular taskbeing done.
In an industrial situation it might be required to mix up asalt solution for pickling olives (for example). To make iteasy and efficient, the instructions might be
�“1 kg of salt to every 10 litres of water�”or some such.
In this case the units of concentration would be kilogramsper litre (kgL-1).
In another situation, it might be convenient to use %m/v.
In Chemistry, it is usually best to measure in molL-1
(�“molarity�”) because this allows easy conversions of mass,volumes of gases and volumes of solutions, when chemicalreactions are involved.
Technique For Making SolutionsOne important laboratory technique is that of making up asolution to a required concentration.
The first step is to calculate the mass of solute required tomake the desired solution, as in Example Problem 2, on theleft of this page.
Once this exact mass is weighed out, the technique is:
Note that to make 500mL of solution you do NOT add500mL of water. You make the volume of the solution upto 500mL... yes, there IS a difference!
Once a solution is prepared this way, other solutions can bemade from it by taking measured quantities, and dilutingthem appropriately.
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Concentration = number of moles (of solute)(of solution) Volume (of solution)
c = n V
Units of measurementc in moles per litre (molL-1)n in moles (mol), and remember that n = m V in litres (L) MM
Example Problem 1If 12.00g of pure solid NaCl was dissolved in water, andmade up to 250.0mL (0.2500 L) of solution, what is themolar concentration (�“molarity�”) of the solution?
Solution:Step 1. Find the number of moles. MM(NaCl) = 58.44g
Diluting to a Desired ConcentrationA common procedure in the Chemistry laboratory is tohave chemical solutions already prepared to a knownconcentration, and dilute them to new concentrations asneeded.
To calculate the new concentration, or to calculate thevolume needed to get a desired concentration, use thefollowing relationship:
Try the Worksheet, at the end of Section
Equipment for Diluting SolutionsYou may have done practical work inthe laboratory to learn how to carryout a dilution.
The required volume (calculated as inExample Problem 2 on the left) ismeasured by pipette and transferred toa volumetric flask.
Pure water is added to the mark.
These are bulbpipettes whichmeasureaccurately a singlevolume e.g. 25.00mL
For odd amounts(like 5.6mL) use agraduated pipette.
The Concentration of Ions in SolutionWhen an ionic compound dissolves in water the ioniclattice disintegrates as the individual ions are hydrated andtaken into the solution. What is the concentration of theindividual ions?
If the compound contains ions in a 1:1 ratio this is a verysimple situation. For example, consider the dissolving ofsalt, sodium chloride:
NaCl(s) Na+(aq) + Cl-(aq)
If the solution has a concentration of (say) 0.5 molL-1, thenthe concentration of the Na+ ions is 0.5 molL-1 and theconcentration of the Cl- ions is 0.5 molL-1 as well.
However, if magnesium chloride (MgCl2)dissolves there are2 chloride ions for every 1 magnesium ion. If theconcentration of the solution was 0.5 molL-1, then theindividual ion concentrations are:
MgCl2(s) Mg+2(aq) + 2Cl-(aq)
0.5 molL-1 0.5 molL-1 1.0 molL-1
In a 0.5 molL-1 solution of aluminium sulfate theconcentrations would be:
Al2(SO4)3(s) 2Al+3(aq) + 3SO4
-2(aq)
0.5 molL-1 1.0 molL-1 1.5 molL-1
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c1V1 = c2V2
(or cV = constant)
c1 = concentration of original solution, in molL-1
V1 = volume of original solution used, in L **c2 = concentration of diluted solution, in molL-1
V2 = volume of diluted solution made, in L **
** It actually doesn�’t matter what units you use, so longas you are consistent throughout the calculation. In theexamples below, volumes are in mL.
Example Problem 1If 25.00mL of a solution of concentration 0.3750molL-1
was diluted to a new volume of 500.0mL, what is theconcentration of the diluted solution?
Solutionc1V1 = c2V2, so c2 = c1V1/V2
= 0.3750 x 25.00/500.0## c2 = 0.01875 molL-1
(1.875 x 10-2)Example Problem 2It is required to make 250.0mL of a solution withconcentration 5.000x10-3 molL-1, from a �“stocksolution�” with concentration 0.2250molL-1. Whatvolume of the stock solution should be measured fordilution?
Solutionc1V1 = c2V2, so V1 = c2V2/c1
= 5.000x10-3 x 250.0/0.2250## V1 = 5.555 mL
(In fact, you would not be able to measure such a precisevolume by pipette. Appropriate answer is really 5.6 mL)
Bulb Pipettes
Something worth knowing:In Chemistry, square brackets around a formula isshorthand for �“molar concentration of...�”e.g. [NaCl] means �“molar concentration of NaCl�”
Mass, Volume & Concentration in Precipitation ReactionsArmed with a knowledge of molarity you can now link calculations involving
concentration of solutions to masses and even gas volume quantities.
Example Problem 115.00mL of 0.3055 molL-1 solution of lead(II) nitratewas treated as follows:An excess of potassium iodide solution was added,causing a precipitate. The solid precipitate was collectedby filtration, dried and then weighed.
What substance, and what mass, was collected?
(Note: an �“excess�” of something means that the quantityadded was more than enough to ensure a completereaction)
SolutionStep 1: use the �“Solubility Rules�” to figure out whatsubstance precipitated, then write a balanced equationfor the reaction.
Pb(NO3)2(aq) + 2KI(aq) PbI2(s) + 2KNO3(aq)
Step 2: find how many moles of Pb(NO3)2 were presentin the 15mL (0.015 L) of solution.
c = n/V, so n = cV = 0.3055 x 0.01500n(Pb(NO3)2) = 4.5825 x 10-3 mol
Step 3: find how many moles of PbI2 were precipitated.
The balanced equation shows the mole ratio is 1:1,so n(PbI2) = 4.5825 x 10-3 mol
Step 4: convert moles to mass. MM(PbI2) = 461.0gn = m/MM, so m = n x MM
= 4.5825x10-3 x 461.0# m(PbI2) = 2.113g
Note:The working above assumes 100% precipitation of thelead ions. Technically, a small fraction of the lead ionswould stay in the solution, so not quite all of it wouldprecipitate. However, the solubility of PbI2 is very low,so for simplicity (K.I.S.S.) we�’re assuming completeprecipitation.
Example Problem 2To measure the concentration of salt in a 40.0mLseawater sample, an excess of silver nitrate solution wasadded to precipitate all the chloride ions. The precipitatewas collected by filtration, dried and weighed. Its masswas 2.76g
a) What substance was precipitated?b) Write a net ionic equation for the precipitation.c) Write a full ionic equation for the reaction.d) Calculate the number of moles of precipitate
collected.e) How many moles of chloride ions must have been in
the seawater sample?f) Calculate the molar concentration of salt in the
seawater.
Solution
a) From �“Solubility Rules�”: Silver chloride, AgCl
b) Ag+(aq) + Cl-(aq) AgCl(s)
c) Ag+(aq)+ NO3
-(aq)+ Na+
(aq)+ Cl-(aq)
AgCl(s)+ NO3-(aq)+ Na+
(aq)
d) n = m/MM MM(AgCl) = 143.35g= 2.76 / 143.35
n(AgCl) = 0.0193 mol
e) Mole ratio in equation is 1:1## n(Cl-) = 0.0193 mol
f) c = n/V= 0.0193 / 0.040 (40mL = 0.040 L)
c(NaCl) = 0.481 molL-1
Example Problem 3 A little revision of Topic 2What volume of hydrogen gas (measured at SLC) couldbe produced from the complete reaction of 50.0mL of1.50 molL-1 hydrochloric acid with magnesium.
SolutionAs usual, start with a balanced equation:Mg(s) + 2HCl(aq) H2(g) + MgCl2(aq)
Moles of HCl present in the solution:c = n/V, so n(HCl) = cV = 1.50 x 0.050 = 0.075 mol
Moles of H2: equation shows mole ratio = 2:1# n(H2) = 0.075/2 = 0.0375 mol
Volume of H2: (remember 1 mole = 24.8 L at SLC)vol(H2) = 0.0375 x 24.8 = 0.930 L (930 mL)
Worksheet next page
Worksheet 3Practice Problems
1. Ionic Equations for DissolvingWrite ionic equations (showing states) for the dissolving ofeach compound in water.a) potassium bromideb) calcium sulfatec) lithium nitrated) magnesium iodidee) aluminium nitratef) ammonium chlorideg) iron(II) nitrateh) copper (II) sulfatei) calcium hydroxide
2. Predicting PrecipitatesUse the �“Solubility Rules�” to predict the result of mixingeach pair of ionic solutions.To answer, write �“No Reaction�”, or name the compoundwhich would form a solid precipitate.
3. Ionic EquationsFor each of the combinations in Q2 which would react toform precipitates:
i) write the full ionic equation, and balance.ii) write the net ionic equation
4. Molarity Calculationsa) What is the �“molarity�” of a solution if:
i) 2.50 mol is dissolved in 0.750 L of solution?ii) 0.025 mol is dissolved in 0.050 L of solution?iii) 0.35 mol is dissolved in 100 mL of solution?iv) 1.2x10-3mol is dissolved in 4.0 L of solution?v) 0.95 mol is dissolved in 200mL of solution?
b) How many moles of solute are ini) 2.00 L of a 0.400 molL-1 solution?ii) 0.450 L of a 1.25 molL-1 solution?iii) 50mL of a 0.025 molL-1 solution?iv) 2.00mL of a 0.0035 molL-1 solution?v) 0.050 L of a 2.25 molL-1 solution?
5. Molarity & Mass Calculationsa) What is the molar concentration of each solution?
Mass of solute dissolved in Volume of Solutioni) 15.80g of potassium nitrate 0.200 Lii) 3.66g of copper(II) sulfate 500mLiii) 127g of sodium chloride 1.50 Liv) 85.6g of lead(II) nitrate 3,000mLv) 2.35g of lithium bromide 250mL
b) What mass of solute is required to make each solution?Solute Concentration (molL-1) Volume
6. Dilution of Solutionsa) What is the concentration of the diluted solution if:i) 25.0 mL of 0.100 MolL-1 solution was diluted to 1.00L?ii) 5.00 mL of 1.25 MolL-1 solution was diluted to 100mL?iii) 2.5 mL of 0.025 MolL-1 solution was diluted to 0.50L?iv) 8.6 mL of 0.500 MolL-1 solution was diluted to 50mL?v) 10.0mL of 5.35x10-3 MolL-1 sol. was diluted to 250mL?
b) In each case, what volume of the �“stock solution�” isneeded to make up the given volume at the requiredconcentration?
7. Mass & Concentration in Precipitation Reactionsa) An excess of sodium sulfate solution was added to25.0mL of a 0.500 molL-1 solution of barium nitrate.Assuming complete precipitation, calculate the mass ofdried precipitate which could be collected.
b) The same reaction as in (a) was used to analyse theconcentration of a solution of sodium sulfate. A 10.0mLsample of the solution was taken, and an excess of bariumnitrate solution was added. The mass of dried precipitatecollected was 1.27g.What was the concentration of the sodium sulfate solution?
c) A precipitate of silver carbonate was collected from50.0mL of 0.500molL-1 solution of silver nitrate, by addingan excess of potassium carbonate solution.
i) Find the mass of dried silver carbonate collected.ii) (Some Revision!) If this silver carbonate was heated
and decomposed, what mass of silver metal, and whatvolume of carbon dioxide gas (at SLC) would be formed?
Temperature, Heat Energy & Heat CapacityWhen heat is added to any substance, what really happensis that the particles (atoms/ions/molecules) move faster. Insolids the particles just vibrate more quickly, in liquids orgases they actually move around faster.
What we measure and understand as �“temperature�” is reallya measurement of the average kinetic energy (movement)of the particles.
Not all particles speed up equally when heat is added:
If you do the same thing to water:
The temperature of the water does not change much whenheat is added.
(Explanation: it�’s those �“sticky�” polar molecules again!Water molecules cling to each other by hydrogen bonding.This means they are hard to accelerate, and it takes moreenergy to make them speed up.)
�“Specific Heat Capacity�” is a measure of how much heatenergy (in joules) is required to change the temperature of1 gram of a substance, by 1oC. The units of Heat Capacityare, therefore, joules per degree per gram (J/oC/g)
Comparison of Some Specific Heat Capacities
Substance Heat Capacity (J/oC/g)
Water 4.18
Typical Metal 0.3 (approx)
Other Liquid SolventsEthanol (alcohol) 2.44Acetone 2.17Petrol (mixture) 2.2 (approx)
Note that water�’s Heat Capacity is much higher thanmost other substances......another of water�’s �“weird�” and unusual properties
Measuring Heat Energy ChangesWhen any substance gains or loses heat, the amount ofheat energy involved depends upon:
�• the amount of substance. i.e the mass.�• the Specific Heat Capacity of that substance.�• the temperature change.
$H = change in heat energy, in joules (J)m = mass of substance, in grams (g)C = Specific Heat Capacity, in J/oC/g$T = temperature change, in degrees celsius (oC)
Notes�• The Greek letter delta ($) means �“change in...�”�• Chemical Data Sheets may give Heat Capacities for
1 kilogram of substance instead of 1 gram.No problem; just divide by 1,000.
�• Why is there a negative sign??For technical reasons (explained later) if thetemperature goes up, the energy change is considerednegative. If temperature drops (negative temp. rise), theenergy change is considered positive.The negative sign in the equation takes care of this.
Example Problem 1How much energy is needed to raise the temperature of50.0g of water by 12.0oC?Specific Heat Capacity of water = 4.18 J/oC/g
Solution$H = - mC$T
= - 50.0 x 4.18 x 12.0= - 2,508 J
(In this non-chemical situation the (-ve) sign can really beignored. The energy required is 2.51 x 103 J (2.51 kJ))
Example Problem 2If 10,000 J of heat energy was added to 100g of ethanol(Specific Heat Capacity = 2.44 J/oC/g) what would bethe temperature rise?
SolutionSince the temperature will rise, technically the energy is anegative quantity, so $H = - 10,000J
$H = - mC$T,so $T = $H/(-m x C)
= -10,000/(-100 x 2.44)= 41.0 oC i.e. Temp. will rise by 41oC
Calorimetryis a technique used to measure the energy change occurringduring chemical processes. The word is derived from the�“calorie�”, a unit for heat energy no longer in use. Theequipment used to make energy measurements is called a�“calorimeter�”.
(Since we now use joules for our energy unit, maybe weshould call it a �“joulemeter�”)
Since many chemical processes occur in water, and becausewater has such a high Specific Heat Capacity (i.e. it canabsorb lots of energy with little temperature change)calorimetry often uses water as the �“working fluid�” ormedium used to absorb the heat energy.
Endothermic and Exothermic Changeswere introduced in Topic 1. Here is a quick reminder:
Exothermic Reactions (�“Exo�”= to go out)are the reactions that produce and release energy.
The amount of energy involved is the �“delta-H�” for theprocess, and is measured per mole of the substance(s)involved. When the chemicals lose energy, the temperaturein a calorimeter rises, because the energy release heats upthe water in the calorimeter. This is why, when thetemperature rises, the energy quantity is considerednegative... the chemicals involved have LOST this amountof heat energy.
Endothermic Reactions (�“Endo�” = to go in)are the reactions that absorb energy... those where you mustsupply energy to make it happen.
The �“delta-H�” for this change is considered positivebecause the chemicals have gained energy. The temp.change is negative, because the calorimeter temp. drops.
Simple Laboratory Calorimeter
Thermometer measurestemperature change
Copper Beakerreaction container
Polystyrene body and lidprevents heat loss/gainwith the surroundings
Practical Work:Heat of Solution
You may have carried out experiments to measure theenergy change that occurs when ionic compoundsdissolve in water.
General Method:Use a calorimeter to measure the temperature change in ameasured mass of water, when a measured mass of a soliddissolves.
You can then calculate:�• the energy change occurring (for the quantities used)and then,�• the energy change per gram of solute.and then,�• the energy change per mole of solute.
Typical Results for dissolving of Potassium hydroxideMass of water placed in calorimeter = 100gMass of potassium hydroxide dissolved = 4.50gInitial temperature of water = 21oCFinal temperature of solution, after dissolving = 28oC
# Temperature change, $T = 7.0oC
Calculations:$H = - mC$T
= - (100 + 4.5) x 4.18 x 7.0= - 3, 058 J for the dissolving of 4.50g
Energy per gram: $H = - 3,058/4.50 = -679 J per gram
Energy per mole: MM(KOH) = 56.1g$H = -679 x 56.1 = - 38,100 J per mole
Limitations of CalorimetryWhen you use a simple calorimeter to measure an energychange in the laboratory, there are a number ofassumptions and approximations involved.
�• It is assumed that the calorimeter itself does not absorb asignificant amount of the heat energy of the reaction.
This source of error is minimized by using a copperreaction vessel, since the very low Specific Heat Capacity ofcopper means it absorbs little energy.
�• It is assumed that there is no heat lost or gained betweenthe calorimeter and the surroundings.
This source of error can be minimized by good heatinsulation of the calorimeter.
�• It is assumed that the Specific Heat Capacity of thesolution reacting in the calorimeter is the same as water.i.e. C = 4.18 J/oC/g
For many solutions this is not quite true, but (generally) theerror this causes is very small.
A serious limitation of many calorimetry experiments inschool laboratories is the poor precision of the usual lab.thermometers. Usually these can only be read to the nearest0.5oC, and if the temperature change is only a few degrees,the % error is huge. Serious calorimetry needsthermometers with a precision of at least 0.1oC.
Water�’s Heat Capacity & Life on EarthThe fact that water has a remarkably high Specific HeatCapacity is of enormous significance to weather, climateand life on Earth.
It means that, on a hot day, the ocean or a lake can absorba large amount of energy from the Sun without muchtemperature change. The air and the land may get very hot,but the water temperature changes very little. In coldweather, the air and land can get really cold, but the waterchanges only a little.
This means that water habitats have very stabletemperatures and do not change much from day to night,or even summer to winter. Aquatic organisms do not needcomplex temperature control mechanisms because theirhabitat remains quite stable.
More importantly, the oceans absorb and transport (viaocean currents) huge quantities of heat from the tropicstowards the poles. This has the effect of cooling thetropical areas and warming the temperate regions, andgenerally evening-out the Earth�’s temperature.
Without water, very little of the Earth would have liveabletemperatures. Without the �“moderating effect�” of water,the tropics would be too hot for life, and the temperateregions would be too cold.
Thermal PollutionSome industries, especially coal-burning or nuclear powerstations produce large amounts of waste heat.
In some places, these plants are situateded beside lakes orthe sea so that the water can be used for cooling theequipment. Typically, lake water is pumped through theequipment, then hot water discharged back into the lake.
This is thermal pollution, and is very destructive toaquatic habitats.
The main problem is a matter of solubility.
Oxygen, and other gases, are sparingly soluble in water.Aquatic organisms are totally dependent on this lowconcentration of dissolved gases for their survival.
The problem is, that the solubility of gases decreases asthe temperature rises. If the water temperature rises by aslittle as 5oC, the dissolved oxygen concentration drops by20% and fish begin to suffocate.
Not only that, but increased temperatures can interferwith the normal breeding cycles and alter the delicatebalance between populations of food plants, diseasemicrobes, parasites, etc. Habitat destroyed!Photo by Daniel West
Worksheet 4
Part A Fill in the blanksTemperature is a measure of the average a)...........................energy of the particles within a substance. When heat isadded, the particles b)............................................. However,some substances require more heat energy than others forthe same temperature change. This difference is measuredby the property of c)...................................................................which has units of d)............................... Water has a verye).................................. (high/low) value.
The amount of heat energy involved in any change is givenby the formula f).......................................
A g).................................... is a device for measuring energychanges. Water is often used as the �“working fluid�” becauseof its high h)................................................... Capacity.
i)...................-thermic changes release energy, so thetemperature in the calorimeter j).................................... Thechemicals in the reaction have k).......................... energy, sothe energy value is considered l)........................... (+/-)
m)...................-thermic changes absorb energy, so thetemperature in the calorimeter n).................................... Thechemicals in the reaction have o).......................... energy, sothe energy value is considered p)........................... (+/-)
The energy change involved in dissolving a solution iscalled the �“q)..............................................................................�”
Calorimetry has a number of limitations and sources oferror;�• the calorimeter itself may r)..................................................This error is minimized by using a reaction container witha very low s).............................................................................�• heat may be lost or gained between the calorimeter andthe t)............................................ This error is minimized byu)............................................ the calorimeter.�• it is assumed that the reacting solution has the samev)....................................................................... as water. This isan approximation, but only causes a w)........................ error.�• Experimental error often comes from the lack ofprecision of the x)............................................
Water�’s very high S.H.C. is of great significance to theEarth�’s y)..................................... and ...................................Water habitats have very z)......................................temperatures, and the ocean currents aa)................................huge amounts of heat, ab)...................................... the tropicsand warming the ac)................................................ regions.
Thermal pollution is the release of ad)....................................into aquatic habitats. This is destructive, mainly because theae)................................ of gases (such as af).............................)becomes much ag)............................................... at highertemperatures.
Part B Practice Problems
1. Simple Heat Calculations.(The +/- signs may be ignored)
a) Calculate the amount of heat energy involved to:
i) heat 50.0g of water from 20oC to 50oC.ii) cool 400g of water from 95oC to 10oC.iii) heat a swimming pool containing 560 tonnes of water(1 tonne = 1x106 gram) from 12oC to 28oC.iv) heat 100g of copper (C = 0.39 J/oC/g) from 10oC to itsmelting point, 1,085oC.v) cool a 100 kg car engine (steel, C = 0.45 J/oC/g) from120oC to 20oC.
b) Calculate the Final Temperature (nearest degree) if:
i) 250g of water at 20oC absorbs 72,000 J of heat.ii) 5.00 kJ of energy was extracted from 80.0g of water at
25oC.iii) 1 L of water (= 1kg mass) at 4oC absorbs 10,000J.iv) a 5.00kg lump of steel at -25oC absorbs 20,000J.v) 20.0g of ethanol (C = 2.44 J/oC/g) at 30oC loses 1.2kJ.
2. Heat of Solution Calculations (+/- sign important!)
a) Find the Molar Heat of Solution if:
i) when 5.85g of ammonium nitrate dissolved in 100mL ofwater (100mL water = 100g) in a calorimeter, thetemperature went from 24oC to 11oC.
ii) 8.42g of sodium carbonate was dissolved in 100mL ofwater in a calorimeter. The temperature increased by 9.5oC.
b) Find the Final Temperature
i) The �“heat of solution�” for sodium hydroxide is listed as$Hsol (NaOH) = -41.6 kJmol-1.
If 10.0g of NaOH was dissolved in enough water to make250g of solution, what would the final temperature be? Theinitial temperature was 18oC.
ii) The �“heat of solution�” for ammonium chloride is listedas $Hsol (NH4Cl) = +15.2 kJmol-1.
If 18.5g of NH4Cl was dissolved in 150mL of water(initially at 22oC) what would be the final temperature?
CONCEPT DIAGRAM (�“Mind Map�”) OF TOPICSome students find that memorizing the OUTLINE of a topic
helps them learn and remember the concepts and important facts.Practise on this blank version.
Practice QuestionsThese are not intended to be "HSC style" questions, but tochallenge your basic knowledge and understanding of thetopic, and remind you of what you NEED to know at theK.I.S.S. Principle level.
When you have confidently mastered this level, it isstrongly recommended you work on questions from pastexam papers.
Part A Multiple Choice
1.The water content of the lithosphere is mainly in the formofA. water vapourB. liquid waterC. solid iceD. water of crystallization in minerals
2.The �“density anomaly�” of water is that:A. the density of ice is higher than liquid water.B. the density of water is exactly 1.00 g/mL.C. ice floats in water.D. water�’s density is extremely high.
3.The compound H2Se can be represented by the Lewis Formula shown.The covalent molecule contains 2 unshared pairs ofelectrons. You would expect its shape to be:A. linearB. tetrahedralC. triangular pyramidD. bent
4.The �“inter-molecular forces�” in water are really:A. hydrogen bondsB. polar covalent bondsC. pure covalent bondsD. ionic attractions
5.�“Hydrogen Bonds�” are likely to occur within substances inwhich hydrogen atoms are bonded to atoms of:A. oxygen, chlorine and carbon.B. nitrogen, oxygen and fluorine.C. sulfur, oxygen and chlorine.D. fluorine, chlorine and bromine.
6.Water tends to form droplets because of its:A. high viscosity.B. high surface tension.C. density anomalyD. high boiling point.
7.If small, non-polar molecules were mixed with water, youwould expect them to:A. dissolve, as separate hydrated ions.B. dissolve, hydrated by hydrogen bonding.C. dissolve, dispersed in water at low concentration.D. not dissolve at all.
8.The correct ionic equation for the dissolving of solidaluminium chloride is:A. AlCl3(s) AlCl3(aq)
B. AlCl3(s) Al+3(aq) + Cl-(aq)
C. AlCl3(s) Al+3(aq) + 3Cl-(aq)
D. AlCl3(s) 3Al+(aq) + Cl-3(aq)
9.If solutions of potassium carbonate and calcium nitratewere mixed together, you would observe:A. no reaction.B. a precipitate of potassium nitrate.C. a precipitate of calcium carbonate.D. a precipitate of potassium calcide.
10.A saturated solution of lithium bromide is in the samebeaker with solid crystals of lithium bromide. Theconcentration of ions in solution does not change overtime because:A. the rates of dissolving and precipitating are the same.B. there can be no further dissolving because the solution
is saturated.C. all chemical processes have ceased at equilibrium.D. the lithium & bromine ions are at equal concentrations.
11.The number of moles of sodium ions in 500mL of a 2.0 molL-1 salt solution, is the same as the number of molesof chloride ions in:A. 500mL of a 2.0molL-1 solution of MgCl2.B. 1 L of a 1.0molL-1 solution of MgCl2.C. 2 L of a 1.0molL-1 solution of MgCl2.D. 250mL of a 2.0molL-1 solution of MgCl2.
12.To make 2.0 L of potassium bromide solution to aconcentration of 0.1 molL-1 would required a mass of KBrclosest to:A. 2 g B. 10g C. 20 g D. 120g
13.In the process of accurately preparing 250mL of a solutionto a specified concentration, you would need to accurately:A. measure 250.0mL by pipette.B. add water to the mark in a 250mL volumetric flask.C. fill a 250mL measuring cylinder.D. use a 250mL graduated beaker.
14.The Specific Heat Capacity values for 4 substances areshown.
Substance S.H.C.(J/oC/g)A 0.95B 2.5C 1.2D 2.1
Which substance (A, B, C or D) would have the largesttemperature rise, if equal amounts of heat energy wereadded to equal masses of each substance?
15.An exothermic process is considered to have a negativevalue for the energy change, because:A. the temperature in the reaction container drops.B. you need to put energy into the system.C. it is the opposite of an endothermic change.D. the chemicals have lost energy in the change.
16.�“Thermal Pollution�” damages ecosystems mainly because:A. living things cannot tolerate the temperature rise.B. less oxygen can dissolve in warmer water.C. the extra heat speeds up plant growth.D. more dirt dissolves in warm water, so erosion increases.
Longer Response QuestionsMark values shown are suggestions only, and are to giveyou an idea of how detailed an answer is appropriate.
17. (6 marks)List 2 important roles of water:
a) in living things.b) as a factor in weather , climate and geography.c) as a resource for human civilization.
18. (4 marks)a) What is meant by saying that water has a �“densityanomaly�”?b) Explain, in terms of bonding and particle arrangements,why water has a density anomaly.
19. (4 marks)a) Draw a Lewis Formula for the water molecule.b) Explain the shape of the molecule
i) a covalent bond.ii) a hydrogen bond.iii) the partial charges (!+ , !-)
on one molecule.b) Explain what is meant by a �“polar covalent bond�”.c) Explain how the presence of hydrogen bonds isresponsible for water�’s relatively high m.p. & b.p.
21. ( marks)a) What is meant by the �“viscosity�” of a liquid?b) Describe a laboratory exercise you have done to comparethe viscosity of 2 or more different liquids.c) How does the viscosity of water compare to that ofother liquids with similar sized molecules?d) Discuss the significance to aquatic animals, of youranswer to part (c).
22. (6 marks)Write ionic equations to describe:
a) the dissolving of calcium chloride in water.b) the precipitation of calcium chloride from water.
(For example, if a solution was evaporated to dryness)c) the situation of a saturated solution of calcium chloride,
in contact with solid calcium chloride.
23. (10 marks)a) Predict what would happen if solutions of potassiumcarbonate and lead(II) nitrate were mixed together.
b) Write a full ionic equation to describe the reaction.
c) The lead(II) nitrate solution had a concentration of 0.500molL-1 and 20.0mL was used. Calculate the mass ofprecipitate formed. (Assume complete precipitation, afteradding an excess of potassium carbonate solution.)
d) Write a symbol equation to describe the decompositionof the dried precipitate, given that a pure metal is formed,and a mixture of 2 gases.
e) Calculate the total gas volume (measured at SLC) formedby the decomposition of the quantity of precipitate formedin part (c).
24. (6 marks)You have been given the task of preparing 500mL of asolution of potassium iodide (KI) with a concentration of0.250 molL-1, from the solid pure chemical.
Describe the steps of the procedure, including the exactmass you would use, and any points of technique to ensureaccuracy.
25. ( marks)Using a polystyrene cup as a simple calorimeter, a studentadded 50.0mL of water and measured its temperature to be18oC. She weighed out 4.27g of lithium hydroxide anddissolved it in the water. The water temperature rose to amaximum of 45oC.
a) Showing all working, calculate the molar Heat ofSolution (including sign) for lithium hydroxide.
b) Later, she looked up a Chemical Data Book and foundthe �“accepted�” value for $Hsol(LiOH). It was a significantlylarger amount of energy than the experimental results gave.Suggest 2 reasons why.
