2.
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
3.
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
4. 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.
7.
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
8. 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
9.
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
10.
11.
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
12.
13.
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
14.
15.
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
16.
17.
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
18.
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
19.
20.
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.
21.
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
22.
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.
23.
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.
27. 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
28.
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.
29.
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.
30.
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
31.
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
32.
(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
33.
(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
34.
Redox reaction is a reaction with
involves both oxidation and reduction. In a nutshell,
OIL RIG happens simultaneously
Redox
35.
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
37.
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
38.
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
39.
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
40. pH 1-6 substances are usually acidic
pH 7 substances are usually neutral
pH 7-14 substances are usually alkaline
41.
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
42.
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
46.
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
47.
Acid Base Reaction
Acid + Base —-> Salt + water
We call this the Neutralization Method.
48.
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.
49.
Acid + Metal Hydroxide —-> Salt +
Water
Hydrochloric Acid + Sodium Hydroxide —-> Water + Sodium
Chloride
Characteristics of the Reaction
Hydroxide starts to disappear
Temperature increases (exothermic reaction)
50.
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
51.
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.
52.
Controlling Acidity in
soil
To reduce the acidity, the soil is treated with a base
like limestone or quicklime or slaked lime
53.
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.
54.
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).
55.
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.)
56.
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)
57.
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
58.
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
59.
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).
60.
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)
61.
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
63. 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).
65. 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.
72. 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.
73.
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
74.
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
75.
76.
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
77. 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
78.
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
79.
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)
80.
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.
81.
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)
82. 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.
83.
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).
84.
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.
85.
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
86.
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.
87.
88.
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
89.
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.
90.
91.
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
92.
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.
93.
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.
