Chemistry review 6 9
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Chemistry review 6 9 Chemistry review 6 9 Presentation Transcript

  • Mary Rodriguez
  •   1 Relate the terms exothermic and endothermic to the temperature changes observed during chemical reactions.  2 Demonstrate understanding that exothermic and endothermic changes relate to the transformation of chemical energy to heat (thermal energy), and vice versa. C6. Energy changes in chemical reactions
  •   During all chemical reactions, an energy change occurs. In the reaction, heat is either released or absorbed. When a reaction releases heat to the surroundings, we call that reaction an Exothermic Reaction. The reaction that absorbs energy from the surroundings are called Endothermic Reactions. Endothermic and Exothermic
  •  Exothermic Reaction  The reactants have more energy than the products here, so a small amount of energy is required to activate the reaction.  Release of heat  Energy needed for the reaction to occur is less than the total energy released.  Extra energy is released, usually in the form of heat.  The release of heat means that an exothermic reaction increases temperature of the surroundings.  Endothermic  Heat absorbs energy from the surroundings.  Temperature of surroundings decreases during an endothermic reaction because energy from surroundings is required to drive the reaction, hence decreasing the temperature of the surroundings.
  • Exothermic
  •  Energy Transformation  In order to actually start a reaction, a certain amount of energy will be provided to the reactants; We often call this the Energy of Activation because this energy is essentially required to start the reaction.  The energy here is used to break the bonds between the molecules of the atoms of the reactants. The bonds then subsequently rearrange and bond again, which releases energy.  However, if the energy provided to activate the energy is less than the energy released when the bonds form together, the reaction gave out more than it took/absorbed, which makes this a exothermic reaction. If the energy given to activate is more than the energy released during the bond formation, the reaction is endothermic.  The total energy change is called enthalpy
  •  1 Describe the effect of concentration, particle size, catalysis and temperature on the speeds of reactions.  2 Describe a practical method for investigating the speed of a reaction involving gas evolution.  5 Describe the application of the above factors to the danger of explosive combustion with fine powders (e.g. flour mills) and gases (e.g. mines).  7 Define catalyst as an agent which increases rate but which remains unchanged.  3 Devise a suitable method for investigating the effect of a given variable on the speed of a reaction.  4 Interpret data obtained from experiments concerned with speed of reaction.  6 Describe and explain the effects of temperature and concentration in terms of collisions between reacting particles (concept of activation energy will not be examined). C7. Chemical reactions
  •   Increasing the surface area will subsequently increase the rate of reaction, as increasing surface area will increase the chances of particles colliding with each other and will hence increase the rate of reaction.  In the exam, they might ask you questions like, ―which reacts faster, magnesium or magnesium powder‖. The obvious answer is the powder because the powder has a much larger surface area, hence increasing the rate of reaction. Surface Area
  •   Increasing the concentration increases the rate of reaction, as there will be collisions per second per unit volume.  The reason for this is increasing the concentration results in there being more particles in each cm3 of space, so there will be more frequent collisions between particles.  As the reaction occurs and the reactants get slowly used up, the concentration of the substance then decreases. This explains for a slower rate of reaction as the reaction proceeds for a period of time. Concentration
  •   Increasing the temperature will increase the rate of reaction. There are two important reasons for this:  Particles will move faster and have more kinetic energy so there will be more collisions per second.  More colliding particles will have the necessary activation energy required, hence allowing more successful collisions.  The second reason is a more important factor in explaining the increased rate of reaction than that of the first. Temperature
  •   Adding a catalyst increases the rate of reaction, but it itself is not used up in the reaction. Catalyst speed up the reaction by lowering the activation energy or providing an alternative pathway for the reacting particles. Catalysts
  •   Speed at which reactants are used up  Speed at which products are formed.  When a gas is produced during a reaction, we can easily measure the reaction by measuring the ―Volume of the gas produced‖. Speed of Reaction
  •   Magnesium + Hydrochloric –> Magnesium Chloride + Hydrogen  We can measure the volume of hydrogen produced. However, to do this, we need to devise an experiment suitable for high school students.  The experiment proposed to measure the volume of gas produced is described below:  Apparatus  Gas syringe  Excess Dilute Hydrochloric Acid  Magnesium  Stopwatch  Conical Flask Example
  •  Steps:  1.Clean Magnesium with sandpaper to ensure any impurities are cleaned off.  2.Put the dilute hydrochloric acid into the flask.  3.Add the magnesium into the flask.  4.Simultaneously, add the stopper + gas syringe.  5.Start stopwatch.  Obviously, you are going to produce hydrogen in this reaction. The hydrogen will show itself in the form of bubbles, and these bubbles will rise up the flask and into the gas syringe, hence pushing the plunger.  6. Measure the volume moved by the plunger every minute.
  •  Concentration  Remember, we did the experiment on Point 2. Repeat the experiment again, however, this time use two different types of concentrations of HCl.  For one experiment use, ―x‖ concentration of HCl  For the second experiment, use ―2x‖ concentration of HCl
  •  Temperature  Again, repeat the experiment but this time with two different types of temperatures of HCl, and compare the differences of Volume of Gas produced.  As you use a higher temperature, we see a steeper graph, hence concluding that a higher temperature leads to a higher rate of reaction.
  •  Surface Area  Repeat the experiment again but this time use two different sizes of magnesium. The magnesium’s you should use are:  1) Normal magnesium chips  2) Magnesium chips of same mass, but smaller pieces.
  •  Catalysts  Again, repeat the experiment two times again, once with a catalyst, and once without.
  •  Interpret Data Time (minutes) Volume of Hydrogen 1 13 2 27 3 34 4 39 5 41 6 41 7 41
  • How can we calculate rate of reaction? We use this simply formula: Rate of Reaction = Volume of Gas Produced / Time. Let’s notice a few things: • The graph is steepest in the beginning. Basically, the rate of reaction is fastest at the beginning. • The graph gets less steep, and as we can see the graph eventually levels off at a plateau, where the Volume of Gas produced doesn’t further increase. • We can’t really calculate the instantaneous change in rate of reaction, but we can calculate the average rate of reaction. Average Rate of Reaction = Total Volume of Hydrogen / Total Time = 41cm3/ 7 minutes = 41/7 cm3 / per minute
  •  Explosive Combustion  We know that increasing the temperature, concentration and surface area of a substance can increase the rate of reaction. However, it’s important that you don’t overdo it. Overdoing any one of these factors can lead to some severe consequences, including explosions. Here are examples of places where adding too much of any of these factors can lead to some serious consequences:  Flour Mills: Flour particles are very tiny, just like all particles. Hence, flour particles tend to have a large surface area. If there is a lot of flour dust in the air, a spark from the machine is enough to cause a reaction between the flour and the spark to form an explosion.  Coal Mines: In a coal mine, you have all sorts of flammable gases in the air. At the right concentration, these gases form an explosive mix with the air, and this is enough to set off an explosion.
  •  Collisions  Temperature: Increasing the temperature of will give the particles more kinetic energy, so the frequency of collisions between particles will increase and the number of successful collisions will also increase.  Concentration: Increasing concentration increases rate of reaction.The reason for this is increasing the concentration results in there being more particles in each cm3 of space, so there will be more frequent collisions between particles.  As the reaction occurs and the reactants get slowly used up, the concentration of the substance then decreases. This explains for a slower rate of reaction as the reaction proceeds for a period of time.
  •   Catalysts increases the rate of reaction, but the catalyst is not used up in the reaction. Basically, you can reuse the catalyst after a reaction because it is unchanged. Catalyst
  •   1 Define oxidation and reduction in terms of oxygen loss/gain, and identify such reactions from given information.  2 Define redox in terms of electron transfer, and identify such reactions from given information. 7.2 Redox
  •   (a) The gain or addition of oxygen by an atom, molecule or ion.  e.g.  (b) The loss or removal of electrons from an atom, ion or molecule.  e.g.  (c). An oxidizing agent is the species that gives the oxygen or removes the electrons. Oxidation
  •   (a) The removal of oxygen in a compound.  E.g.  (the ―O‖ is lost. )   (b) The gain or addition of electrons to an atom, molecule or ion  E.g.   (c) A reducing agent is the species that removes the oxygen and ―donates‖ the electrons.  An easy way to memorize the electron part of oxidization and reduction is : OILRIG  Oxidation Is Loss, Reduction Is Gain. (In terms of electrons) Reduction
  •   Redox reaction is a reaction with involves both oxidation and reduction. In a nutshell, OIL RIG happens simultaneously Redox
  •   Copper (II) Oxide + Hydrogen –> Copper + Water  Let’s see what’s going on here and why this is a redox reaction.  CuO becomes Cu. Is oxygen gained or lost? Lost of course! And we learned, lost of oxygen is reduction which is precisely what’s going on here. CuO is reduced to Cu  Now, let’s look at H2. H2 becomes H2O. Oxygen is gained here, so H2 is oxidized into H2O.  As we can see here, both oxidation (from the H2) and reduction (from the CuO) is taking place. That’s why a redox reaction is occurring Example
  •   iron(III) oxide + carbon monoxide –> iron + carbon dioxide Trial
  •   The iron(III) oxide is reduced to iron, the carbon monoxide is oxidised to carbon dioxide  Since both oxidation and reduction occur here, this is an example of a redox reaction. Answers
  •   1 Describe neutrality and relative acidity and alkalinity in terms of pH (whole numbers only) measured using full-range indicator and litmus.  2 Describe the characteristic reactions between acids and metals, bases (including alkalis) and carbonates.  3 Describe and explain the importance of controlling acidity in the environment (air, water and soil). C8. Acids, bases and salts
  •   All substances are divided into three categories:  Acidic  Alkaline  Neutral  Measured by measuring the pH of the substance. What the pH is that its simply measure of the Hydrogen ion concentration in a substance. However, calculations of that is beyond the scope of the IGCSE Science – if you do, however, want to get a feel of pH calculations, you can visit here.  We measure pH using the pH scale. Describing with pH
  • pH 1-6 substances are usually acidic pH 7 substances are usually neutral pH 7-14 substances are usually alkaline
  •   This is a substance that changes color when it is added to another substance. What color it changes to depends on the pH of the substance. Universal Indicator
  •   This is an indicator also used to test for acidity, neutrality or alkalinity in a substance.  We use something called litmus paper to test for this.  If we want to test for acidity, we use Blue Litmus Paper  If we want to test for alkalinity, we use Red Litmus Paper  The following results are:  Acids: Turn blue litmus paper red.  Alkalines/Bases: Turn red litmus paper blue.  Neutral: No color change. Litmus Paper
  •  Metal Acid Reaction  Metal + Acid —-> Salt + Hydrogen  We call this the ―Displacement‖ method.  Characteristics of the reaction  Bubbles are given out  Temperature rises (the reaction is exothermic, heat is released)  Metal disappears
  •  Acid Base Reaction  Acid + Base —-> Salt + water  We call this the Neutralization Method.
  •  Acid + Metal Oxide —-> Salt + Water Copper Oxide + Sulfuric Acid —-> Copper Sulfate + Water  Here, the Copper merges with Sulfuric acid to make Copper sulfate. If you have iron oxide, nothing will change, the iron will merge with the sulfuric acid to make copper sulfate.  Characteristics of the Reaction  Amount of metal oxide decreases  Temperature increases (exothermic reaction)  Solution changes color.
  •  Acid + Metal Hydroxide —-> Salt + Water  Hydrochloric Acid + Sodium Hydroxide —-> Water + Sodium Chloride  Characteristics of the Reaction  Hydroxide starts to disappear  Temperature increases (exothermic reaction)
  •  Acid + Metal Carbonate —-> Salt + Water + Carbon Dioxide  Sulfuric Acid (Acid) + Copper Carbonate (Carbonate) —-> Copper sulfate (salt) + Water + Carbon Dioxide  Characteristics of reaction  Metal carbonate starts to disappear  Temperature rises (exothermic reaction)  Color Change
  •  Controlling acidity (air,water,soil)  Most crops grow best when the pH of the soil is near 7. If soil is too acidic or too alkaline, crops grow badly or not at all.  Usually acidity is the problem. Why? Because of a lot of vegetation rotting in it or because too much fertilizer was used in the past.  Affects of lower pH  Lack of nutrients  Poor growth of crops  May pass onto rivers, damaging the eco-system within it.
  •  Controlling Acidity in soil  To reduce the acidity, the soil is treated with a base like limestone or quicklime or slaked lime
  •  TYPES OF OXIDES  Acidic and Basic Oxides  􀀁 The oxides that one uses to form acids  and bases in aqueous solution often  have reactivity that reflects their acidic  or basic character.  􀀁 Examples: Li2O, CaO, and BaO react  with water to form basic solutions and  can react with acids directly to form  salts. Likewise, SO3, CO2, and N2O5  form acidic aqueous solutions and can  react directly with bases to give salts.
  •   Oxides as Acid and Basic Anhydrides  Basic Oxides (usually ―ionic‖)  CaO + 2H2O ––> Ca2+ + 2OH–,  moderately strong base  [O2–] + H2O ––> 2OH– K > 1022  Alkali metal and alkaline earth oxides are  basic (dissolve in acid).
  •   Acidic Oxides (Acid Anhydrides)  element-oxygen (E–O) bond not broken on  dissolution  either  an E – O – E group is hydrolyzed by water  or  water is added across a double bond  Acidic Oxides not soluble in water will dissolve in basic  aqueous solutions to produce salts  e.g. As2O3 + 2NaOH(aq) ––> 2NaH2AsO3  (Often seen for anhydrides of weaker acids.)
  •   Amphoteric Oxides  Dissolve in acids or bases - if strong enough.  E.g., BeO, SnO, certain forms of Al2O3  In strong acids: ZnO + 2HCl(aq) ––> ZnCl2(aq)  ZnO + 2HNO3(aq) ––> Zn(OH2)6  2+ + NO3  -  In strong base: ZnO + 2NaOH(aq) ––>  2Na+(aq) + [Zn(OH4)]2– (aq)
  •   Lux-Flood Concept (Oxide Solids)  􀀁 Acid: Oxide ion acceptor  􀀁 Base: Oxide ion donor  – A generalization that includes reactions  between solids when water never gets  involved. E.g.,  CaO + SiO2 􀀁 CaSiO3  3 Na2O + P2O5 􀀁 2 Na3PO4  NaOH + CO2 􀀁 NaHCO3
  •   Other Oxides  Many oxides (particularly of the transition  metals) are difficult to classify as acidic or  basic because redox chemistry is more  important.  e.g. MnO2 + 4HI (aq. conc.) 􀀁  Mn2+(aq) + I2 + 2H2O + 2 I
  •  Highly charged cations with small radii  make for stronger acids:  [Fe(OH2)6]2+ fairly weak, [Fe(OH2)6]3+ is much  stronger. r(Fe3+) < r(Fe2+), the smaller, more  highly charged (more polarizing) cation  withdraws more e– density from coordinated  water.  More than size is involved:  r(Al3+) < r(Fe3+) ionic radii, but [Fe(OH2)6]3+  is stronger than [Al(OH2)6]3+  FeIII–O bonding probably more covalent  (smaller electronegativity diff. than Al/O).
  •   Transition Metals in High Ox. States: Acidic  Metals in very high oxidation states form strong,  largely covalent, bonds with oxygen  ––> weakens O-H bonds!  e.g. CrO 4  2– weak conjugate base of chromic acid  e.g. MnO 4  – very weak conjugate base of  permanganic acid (both are powerful oxidants)
  •  Examples of Acids From Solvolyzed Metals  Aqua Acids (solvolysis) Al3+ solutions are acidic  AlCl3 (s) + H2O → Al3+(aq) + Cl–(aq)  (w/sm. amts of water, HCl gas is evolved):  AlCl3 + H 2 O → ―Al(OH) 3‖ + 3HCl  􀀁 AlCl 2(OH)·nH 2 O complex  AlCl(OH) 2 + mH 2 O
  •  Preparation of Salts
  • Neutralization 1. A known volume of acid is pipetted into a conical flask and universal indicator added. The acid is titrated with the alkali from the burette. 2. The acid is added until the indicator turns green, pH 7 neutral. This means all the acid has been neutralized to form the salt 3. The volume of alkali needed for neutralization is then noted, this is called the endpoint volume. (1)-(3) are repeated with both known volumes mixed together BUT without the contaminating universal indicator. 4. The solution is transferred to an evaporating dish and heated to partially evaporate the water causing crystallization or can be left to slowly evaporate - which tends to give bigger and better crystals. 5. The residual liquid can be decanted away and the crystals can be carefully collected and dried by 'dabbing' with a filter paper OR the crystals can be collected by filtration (below) and dried (as above).
  • Examples: NaOH(aq) + HCl(aq) => NaCl(aq) + H2O(l) NaOH(aq) + HNO3(aq) => NaNO3(aq) + 2H2O(l)
  • Acid and Metal(Insoluble Base) Reaction 1. The required volume of acid is measured out into the beaker with a measuring cylinder. The insoluble metal, oxide, hydroxide or carbonate is weighed out and the solid added in small portions to the acid in the beaker with stirring. 2. The mixture may be heated to speed up the reaction. When no more of the solid dissolves it means ALL the acid is neutralized and there should be a little excess solid. 3. The hot solution (with care!) is filtered to remove the excess solid metal/oxide/carbonate, into an evaporating dish. 4. The hot solution is left to cool and crystallize. Then collect and dry the crystals with a filter paper.
  • Example: CuO(s) + H2SO4(aq) => CuSO4(aq) + H2O(l) CuCO3(s) + H2SO4(aq) ==> CuSO4(aq) + H2O(l) + CO2 (g)
  •  Preparing Insoluble Salt
  •  Making a salt by direct combination of elements
  •  SEPERATION  structure/revise-it/separating-mixtures Filtering, Evaporation, Centrifuging, Crystallizing
  •  PURIFICATION  mdMix2.htm
  • 8.4 Identification of ions and gases  1 Use the following tests to identify: aqueous cations:  ammonium, copper(II), iron(II), iron(III) and zinc by means of aqueous sodium hydroxide and aqueous ammonia as appropriate. (Formulae of complex ions are not required.)  anions:  carbonate by means of dilute acid and then limewater,  chloride by means of aqueous silver nitrate under acidic conditions,  nitrate by reduction with aluminium,  sulfate by means of aqueous barium ions under acidic conditions,  gases:  ammonia by means of damp red litmus paper,  carbon dioxide by means of limewater,  chlorine by means of damp litmus paper,  hydrogen by means of a lighted splint,  oxygen by means of a glowing splint.
  •   1 Describe the way the Periodic Table classifies elements in order of proton number.  2 Use the Periodic Table to predict properties of elements by means of groups and periods. C9. The Periodic Table
  •   1 Describe the change from metallic to non-metallic character across a period.  2 Describe the relationship between Group number, number of outer-shell (valency) electrons and metallic/non-metallic character. 9.1 Periodic trends
  •  Trends  Elements on the left, in Group 1, are all metallic.  Elements in Group 2 are also metallic, but their metallic properties are less apparent than the elements in Group 1.E.g. They are less reactive.  As you go across the group, elements slowly become less metallic, and elements in Group 4 become non-metals. However, they are still generally in the solid form.  As you progress group 6,7,8 elements tend to be in the gaseous form.  The group number is closely related to the number of out- shell valency electrons.  E.g. Sodium is in Group 1 and has 1 outer electron, chlorine is in Group 7 and has 7 outer electrons.  In Group 7, as you go down the group, the substances progress from Gas to solid. e.g. Chlorine is a gas whilst iodine is a solid.  In Group 1 and 2 and possibly 3, even as you down down the group, they are mostly metals, however melting points and boiling points tend to decrease for these substances
  •  1 Describe lithium, sodium and potassium in Group I as a collection of relatively soft metals showing a trend in melting point and reaction with water.  3 Describe the trends in properties of chlorine, bromine and iodine in Group VII including colour, physical state and reactions with other halide ions.  2 Predict the properties of other elements in Group I, given data where appropriate.  4 Predict the properties of other elements in Group VII, given data where appropriate. 9.2 Group properties
  •   1 Describe the transition elements as a collection of metals having high densities, high melting points and forming coloured compounds, and which, as elements and compounds, often act as catalysts. 9.3 Transition elements
  •  Group 1  Melting Points Lithium: 80.5°C (453K) Sodium: 97.8°C (370K) Potassium: 63.38°C (336K) Rubidium: 39.31°C (312K) Caesium: 28.44°C (301K) Francium: 27°C (300.15K)  From here, we can see that as you go down the elements in Group 1, you generally see a decline in melting points.  Boiling points Lithium: 1342°C (1615K) Sodium: 883°C (1156K) Potassium: 759°C (1032K) Rubidium: 688°C (961K) Caesium: 671°C (944K) Francium: 677°C (950.15K)
  •  Lithium, Sodium, Potassium  Lithium floats and then fizzes  Sodium shoots across the water.  The potassium melts with the heat of the reaction, and then the hydrogen catches fire.  Sodium + Water —> Sodium Hydroxide + Hydrogen.  Reactivity Increases as you go down group 1  Alkali’s are known to be very soft, and its softness increases as you go down Group 1 as well.
  •  HALOGENS  Group VII consists mainly of non-metals.  The elements in group VII are:  Fluorine  Chlorine  Bromine  Iodine  characteristics of Halogens that you should be aware of:  Form colored gases: This is quite evident as fluorine is a pale yellow gas and chlorine is a green gas.  Poisonous: Chlorine was typically used in World War I as a poison to kill the enemies. Inhaling chlorine gas is a one-way ticket to the graveyard.  Form Diatomic Molecules: This is basically a molecule that contains two atoms. You will learn more about that in Covalent Bonding (Unit 3.5)
  • Halogen At room temperature: Boiling Points /degrees Fluorine Yellow gas -188 Chlorine Green Gas -35 Bromine Red Liquid 59 Iodine Black Solid 184 •Boiling point increases as we go down the group. This is mostly due to an increased mass which leads to stronger Van Der Waal’s Forces, but this is moving onto IB Chemistry. •Color gets deeper. •Density Increases.
  •  Halogens  Reactivity decreases as you go down the group.  The Halogens often react with metals to form compounds called Halides.  Additionally, as we move down, the elements usually experience a change in state from gas to liquid to solid. Interestingly, the only liquid is bromine at room temperature (25 degrees).
  •  Reactions with other halides  When you add Chlorine water to a solution of Potassium Bromide, the solution turns orange. This reaction is taking place:  Cl2 (aq) + 2KBr (aq) —-> 2KCl (aq) + Br2 (aq)  Chlorine + Potassium Bromide —> Potassium Chloride + Bromine  Bromine has been, what we called, Displaced for a more reactive halogen. It’s a bit like how girlfriends and boyfriends function nowadays. Girl A is with Boy A. The more attractive male, Boy B comes along and displaces Boy A.  What this means is that when a halogen reacts with a salt such as Potassium Bromide (one that contains a -ide), the more reactive halogen will take the place of the less reactive halogen, as we clearly saw in the example just now. Chlorine is definitely more reactive than Bromine, as it is higher up in the Group and reactivity decreases down a group, so it displaces the Bromide in the salt to form Potassium Chloride.
  •  Predict the properties of other elements in Group VII  The trend will continue here and:  Melting and Boiling Points will continue to increase  Becomes darker  Density Increases  Heavier  Example  Astatine  This is one of the rarest elements there are as it decays away extremely rapidly.  However, it is predicted to be the heaviest known element in Group VII.  Highest melting and boiling points of known elements in Group VII
  •  Transition Elements  The elements in the middle section of the Periodic Table are the transition elements. They're all metals with typical metallic properties eg conducting heat and electricity. They often form coloured compounds.  Transition metal carbonates undergo thermal decomposition - a reaction in which a substance is broken down into at least two other substances by heat.  Transition metal hydroxides are insoluble in water. They can be precipitated out of a transitional metal compound solution using sodium hydroxide solution.
  •  Transition Elements  The transition elements are those in the middle section of the Periodic Table.  All the transition elements are metals and so they have typical metallic properties - they conduct heat and electricity, they are malleable and ductile and they form positive ions when they react with non-metals.  The compounds of transition metals are often coloured.  Copper compounds are blue  Iron(II) compounds are light green  Iron(III) compounds are orange/brown  Iron is a catalyst in the Haber process  Nickel is a catalyst used in the manufacture of margarine
  •  Thermal Decomposition  A reaction in which a substance is broken down into at least two other substances by heat is called thermal decomposition.  Transition metal carbonates often undergo thermal decomposition.  The thermal decomposition of copper(II) carbonate is easily demonstrated  If limewater is shaken with a sample of the gas produced, the limewater turns milky. This shows that the gas is carbon dioxide. Notice that the solid in the test tube changes colour as the copper carbonate breaks down to copper oxide and carbon dioxide.
  •   1 Describe the noble gases as being unreactive.  2 Describe the uses of the noble gases in providing an inert atmosphere, i.e. argon in lamps, helium for filling balloons. 9.4 Noble gases
  •  Noble Gases  The noble gases are helium, neon, argon, krypton, xenon and radon.  Radon is radioactive. Radon-220 from rocks is a health hazard.  The noble gases are in group 0 on the right of the periodic table.
  •  Properties  1. The noble gases are almost entirely unreactive because they all have a full outer shell of electrons and are not interested in reacting with other substances. Substances that are very unreactive are called chemically inert.  2. The noble gases are all colourless monatomic gases. Monatomic means that they exist as single atoms. The forces between the atoms are very weak (and so they are gases).  What are the Group Trends for the Noble Gases?  Going down group 0 from helium to radon, the noble gases  1. Have a higher density.  2. Have higher melting points and boiling points because the atoms become heavier (bigger) and require more energy to melt or boil.
  •  USES  tm