The Ultimate IGCSE Guide to Chemistry

7/28/2019 The Ultimate IGCSE Guide to Chemistry

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IGCSE ChemistryIGCSE ChemistryFrom the Edexcel IGCSE 2009 Syllabus including triple science

statements

CGPwned

When Chemistry and CGP books get pwned


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Unit1:ThePeriodicTable

Theperiodictablecontains aboutahundred or soelement

s thathavebeencurrentlydiscovered.Therowsareknownasperiodsandelements of thesameperiodhavethesamenumberofelectronshells.Thecolumns areknown

asgroupsandelements of thesamegrouphavethesamenumber


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of electrons on their outer shell. Group one has one outer electron,

An element is a substance that cannot be broken down into anything simpler. KCl for example (potassium chloride) isNOT an element because it can be broken down into K (potassium) and Cl (chlorine). The potassium and chlorine arethe elements.

A compound is two or more elements chemically bonded together. An example would be KCl (potassium chloride),

which consists of the elements potassium and chloride chemically bonded together.

Atoms are the building blocks of substances.

Molecules are two or more atoms bonded together. It doesnt have to be a compound. Elements such as O2 and Br2are diatomic molecules they exist in pairs.

Atomic Structure

Atoms are made up of protons, neutrons and electrons.Protons are positively charged. Electrons are negativelycharged. Neutrons dont have a charge.

An atom consists of a nucleus, which contains protons andneutrons; and some electron shells which surround the nucleus

and contain electrons. The neutrons however, are different.The number of protons and the number of neutrons add up tomake the mass number of an element.

Understanding the Lack of Reactivity in Noble Gases (Group 0)

Noble gases have eight electrons on their outer shell, therefore, there is no need for them to gain or lose electrons.Basically they have a full outer shell so they dont need to react. This is what makes them so unreactive.

How to Read Each Square on The Periodic Table

You probably already know that the periodic table is made up oflots and lots of squares, each containing an element andinformation about it.

Anyways we already know what the atomic mass number is (thenumber of protons + neutrons). It says 12.011 here but this isprobably because this picture came from some super complicatedperiodic table. In IGCSE level however, the atomic mass should read 12. Anyways, the atomic number is the numberof protons (and electrons), so to find the number of neutrons, if asked to, simply subtract the atomic mass by theatomic number.

Example: Calculate the number of neutrons Carbon has.

The answer: 12 6 = 6 neutrons

The Arrangement of Electrons

Atoms are surrounded by electron shells which contain electrons.But the arrangement is the same for ALL the elements, not matterhow different they are.

Each shell can only hold a certain number of electrons. The very firstshell can hold only two electrons. The second shell can hold eight.The third sometimes appear full with eight but can expand to a totalof eighteen. However, this is beyond GCSE level, and for now, theshells only hold eight.


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So how do you find the electron configuration? Well lets use potassium (K) as an example.Look up the atomic number of potassium. It should say 19. This tells you the number of protons, which is equal to thenumber of electrons so we can use that.

Arrange the electrons in shells, always filling up the inner shell before you go to the outer one. Remember the first,innermost shell can only take 2 electrons, the second one can take 8, and the third one, 8. You will find that you haveone electron left. That goes on the fourth shell.

Your electron configuration should look like this: 2, 8, 8, 1.

Example: Work out the electron configuration of chlorine.

Chlorine has an atomic number of 17 so 17 electrons.17 2 (as the innermost shell only holds two electrons) = 1515 8 (as the second shell only holds eight electrons) = 7 (This number is the number of electrons Chlorinehas on its outer shell).7 electrons does not fill up the third shell so we are left with the configuration: 2, 8, 7.

Isotopes

The number of neutrons in an atom can vary slightly. For example, there are three kinds of carbon atom, called

carbon-12, carbon-13 and carbon-14. They all have the same number of protons, but the number of neutrons vary.These different atoms of carbon are called isotopes. Isotopes are atoms that have the same atomic number, butdifferent mass numbers. They have the same number of protons, but different numbers of neutrons. The fact thatthey have varying numbers of neutrons makes no difference whatsoever to their chemical reactions. The chemicalproperties are governed by the number and arrangement of the electrons.

Calculating Relative Atomic Mass (R.A.M.)

Lets start this off with an example!

Example: Naturally occurring silver is 51.84% silver-107 and 48.16% silver-109. Calculate the relative atomic mass ofsilver.

r.a.m. (Ag) = (51.84/100 x 107) + (48.16/100 x 109)= 55.469 + 52.494=107.96

Now what did we do there? Well I simply calculated 51.84% of 107 (of silver) and 48.16% of 109 (of silver), andadded the two answers! What we end up with is 107.96. Round that up to a whole number and the average relativeatomic mass of silver is 108.

Calculating the Abundance (percentage) of an Isotope

Example: Copper consists of two isotopes, copper-63 and copper-65. Its relative atomic mass is 63.62. Find the

abundance of each isotope.

Let y/100 = abundance of copper-63Let (100-y)/100 = abundance of copper-65

63.62 = (y/100 x 63) + [(100-y)/100 x 65]63.62 = 63y +6500 65y-2y = -135y = 69

Abundance of copper-63 = 69%


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Abundance of copper-65 = 100 69 = 31%

About Metals and Non-Metals

The IGCSE spec. states you have to recall the positions of metals and non-metals on the periodic table. Thats easy! Its on page two. Have a look. Itscolour-coded.

Anyways, this section covers 2.2, 2.3 and 2.5.

Metals

Metals tend to be shiny. They tend to have high melting and boiling pointsbecause of powerful attractions. Metals conduct heat and electricity because delocalized electrons are free to movethroughout the structure. Metals are usually easy to shape due to their regular packed molecules. Metals react withwater to form bases, and their oxides are also bases. They are good reducing agents because they lose electron.

Non-Metals

Non-metals tend to be brittle. They are poor conductors of heat and electricity. They form acidic oxides and are goodoxydising agents because they gain electrons.

Aluminium Oxide

Aluminium oxide is amphoteric. It can neutralize both an acid and a base.

Reaction with acids

Aluminium oxide contains oxide ions and so reacts with acids in the same way as sodium or magnesium oxides. Thatmeans, for example, that aluminium oxide will react with hot dilute hydrochloric acid to give aluminium chloridesolution.

In this (and similar reactions with other acids), aluminium oxide is showing the basic side of its amphoteric nature.

Reaction with bases

Aluminium oxide has also got an acidic side to its nature, and it shows this by reacting with bases such as sodiumhydroxide solution. Various aluminates are formed - compounds where the aluminium is found in the negative ion.This is possible because aluminium has the ability to form covalent bonds with oxygen.

Group 1: The Alkali Metals

Alkali metals are metals that are part of group one. They are extremely reactive metals, and reactivityincreases DOWNWARDS in other words, lithium is the least reactive and francium.

Some Basic Physical Properties

Metal Melting Point ( C) Boiling Point ( C) Density (g/cm )

Lithium 181 1342 0.53


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Sodium 98 883 0.97

Potassium 63 760 0.86

Rubidium 39 686 1.53

Francium 29 669 1.88

You can see that as reactivity increases, the melting and boiling points decreases; however, density increases. Thesepoints are very low for metals. Remember that potassium, sodium and lithium would float on water due to their

densities. But why are they so reactive? Well they only have one electron to lose!

The metals are also very soft and easy to cut, becoming softer as you go down the group. They are shiny and silverwhen cut, but tarnish within seconds on exposure to air.

Storage and Handling

All these metals are extremely reactive. Anyways the metals will quickly react with air to form oxides, and reactbetween rapidly and violently with water to form strongly alkaline solutions of metal hydroxides.

To stop them reacting with oxygen or water vapour in the air, lithium, sodium and potassium are stored under oil.Rubidium and caesium are so reactive that they have to be stored in sealed glass tubes to stop any possibility ofoxygen getting at them.

Great care must be taken not to touch any of these metals with bare fingers. There could be enough sweat on yourskin to give a reaction producing lots of heat and a very corrosive metal hydroxide.

Reactions with Water

All these metals react with water to produce a metal hydroxide and hydrogen.

Metal + Water Metal Hydroxide + Hydrogen

All the hydroxides are bases and turn pH paper purple.

With Sodium

The sodium floats because it is less dense than water. It melts because its melting point is low and a lot of heat isproduced by the reaction. Observations would be that the sodium would turn into a ball and whiz around the surface

of the water. It may form a white trail which is sodium hydroxide. This dissolves to make a strongly alkaline solutionwith the water. When lit, it produces a yellow flame.

With Lithium

The reaction is very similarto sodiums reaction, except it is slower. The lithium does not melt due to its highermelting point. When lit, it produces a red flame.

With Potassium

Potassiums reaction is faster than sodiums. Enough heat is produced to ignite the hydrogen, which burns with a lilacflame. The reaction often ends with the potassium spitting around.

With Rubidium and Caesium The Two Baddies

The reaction is so violent it can be explosive. When lit, Rubidium forms a red flame and Caesium forms a blue flame.

Explaining the Increase in Reactivity

The differences between reactions depend in part on how easily the outerelectron of the metal is lost in each case. That depends on how strongly itis attracted to the nucleus. The more electron shells an atom has, the lesspowerful the attraction forces are. For example, Lithium is a lot lessreactive than Potassium. This is because there are less shells which shieldthe full attraction of the nucleus from the This makes the electron harder to


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lose. However, potassium has a lot more electron shells which shield the outer electron from the nucleus. Thisweakens the attraction in compared to lithium, and therefore, the electron is easier to lose.

Compounds of Alkali Metals

All group one metal ions are colourless. That means that their compounds will be colourless or white unless they arecombined with a coloured negative ion (remember metals would become positive ions because they lose electrons,

whereas, most non-metals gain electrons). Potassium dichromate is orange, for example, because the dichromateion is orange. Group one compounds are typical ionic solids and are mostly soluble in water.

Alkali Metals: Quick Notes

Group One so +1 charge One electron on outer shell Reactivity increases downwards Density increases downwards Melting and Boiling points both decrease downwards Very soft and tarnish quickly in air Li, Na and K are stored under oil, whilst Rb and Cs are stored in

sealed glass tubes Reacts with air to form oxides Reacts with water to form alkaline hydroxides, which turns pH paper

purple Positive ions are formed and they are colourless Flame Colours: lithium, red; sodium, yellow; potassium, lilac;

rubidium, red; caesium, blue. Forget about Francium you dont need to know much about it.

Group 2: Alkali Earth Metals

Alkali earth metals belong to Group two. They are beryllium, magnesium,calcium, strontium, barium and radium. These metals are harder thanthose in group one. They are silvery grey in colour. They tarnish quickly,however they dont just disappear into thin air because the oxides themetals form when reacting with air would form an outer coat that protectsthe metal from the air. They are good conductors of heat and electricity.

They burn in oxygen to form white oxides. They react with water to formhydroxides and hydrogen, but the reaction is a lot less than that of groupone. Also, reactivity increases down the group.

Flame Colours

-Calcium brick red-Strontium -crimson-Barium -apple-green

Well thats it for Group two!

Alkali Earth Metals: Quick Notes

Harder than group one metals Two electrons on outer shell (2+ charge) Form white oxides Forms hydroxides and hydrogen when reacting with water. Reaction is less vigorous than that of group one Reaction increases downwards Silvery-Grey Flame Colours: calcium, brick red; strontium, crimson; barium, apple-green.

Element State Colours


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Group 7: The Halogens

Halogens are group seven elements. Their elements arediatomic molecules. They exist in pairs, such as F2 andCl2. These two elements are gases, bromine is a liquidand iodine is a solid. Astatine is radioactive.

These vapours and gases are poisonous. All these elements need to be handled in a fume cupboard.

Reactions with Hydrogen

The halogens react with hydrogen to formhydrogen halides such as hydrogen fluorideand hydrogen chloride. These are all steamy,acidic and poisonous gases. They are verysoluble in water, reacting with it to producesolutions of acids. However as a gas, it is NOT

an acid.

Reaction Between Sodium and ChlorideSodium burns in chlorine to produce the white solid sodium chloride or salt!

2Na(s) + Cl2(g) 2NaCl(s)

In this reaction, sodium has been oxidized since it has lost electrons. Chlorine has been reduced.

Displacement Reactions with Halogens

Finding the reactivity of halogens aredone by reacting the elements withpotassium halides. Colour changewill indicate a reaction.

Note: Colour changes are due to theelement being displaced. Forexample, the colour change from yellow to brown when chlorine reacted with potassium bromide was dueto the fact that the bromine was displace. It was the brown of the Bromine that turned the solution brown.

Potassium is only a spectator ion. It does not change.

But now we have a problem. To distinguish whether bromine or iodine has been displaced is difficult, asboth elements produce very similar shades of brown. What do we do? We add an organic solvent suchas Volasil. When Volasil is added, the iodine turns pink while the bromine stays brown. Pretty neat huh?

These reactions are known as redox reactions, where oxidation and reduction are occurring (not just oneof them).

Explaining the Trend in Reactivity of Halogens

As you go down the group, the oxidizing ability of the halogens falls due to the decreasing reactivity.When a halogen oxidizes something, it does so by removing electrons from it. Chlorine is a strongoxidizing agent because its atoms readily attract an extra electron to make chloride ions. Bromine is lesssuccessful. Why? This relates to electron shells again. In Chlorine, there are three shells which shield thenucleus attraction force from attracting another electron to gain a full outer electron shell. Brominehowever, has a lot more shells to shield the attraction, therefore, the force is much weaker.

Halogens: Quick Notes

Flourine Gas Yellow

Chlorine Gas Green

Bromine Liquid Orange Brown vapour

Iodine Solid Dark Grey Purple vapour

Halogen Reaction with Hydrogen

Flourine Violent explosion, even in the cold and dark

Chlorine Violent explosion if exposed to a flame or sunlight

Bromine Mild explosion if a bromine vapour/hydrogenmixture is exposed to a flame

Iodine Partial reaction to from hydrogen iodide if vapouris heated continuously with hydrogen

Observations PotassiumChloride

PotassiumBromide

Potassium Iodide

Chlorine Yellow to Brown Yellow to Brown

Bromine Brown Brown to Dark Brown

Iodine Brown Brown


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Diatomic molecules Seven electrons on outer shell Highly reactive only need one electron to fill outer shell Form hydrogen halides when reacting with hydrogen Reaction increases as you go up the group Halogens can displace each other

Volasil turns iodine pink

The Difference Between Hydrogen Chloride and Hydrochloric Acid

Hydrochloric acid is basically a solution of hydrogen chloride gas in water.

The Bronsted-Lowry Theory

Bronsted and Lowry defined acids and bases as the following:

-An acid donates a proton.-A base accepts a proton.

How is this related? Well, when hydrogen loses its only electron, it becomes a hydrogen ion (H+). In other words, it is

also a proton, because it has lost all of its electrons (it only has one remember?).

When hydrogen chloride dissolves in water, a proton (the hydrogen ion) is transferred to the water. This gives us theequation:

H2O(l) + HCl(g) H3O+(aq) + Cl

-(aq)

The H3O+

ion is called a hydroxonium ion. Wenormally write it as H

+(aq). You can think of it as a

hydrogen ion riding on a water molecule.

So in this example, HCl is an acid because itdonates a proton (the hydrogen ion) to water.

So the real differences? Hydrogen chloride is NOT an acid and is a gas. Hydrochloric acid is an aqueous solution ofhydrogen chloride.

Hydrogen Chloride and Methylbenzene

Explaining Water Being a Polar Molecule

Water is a polar molecule. Electronsin water are attracted towards theoxygen end of the bond, whichleaves it slightly negative. Thisleaves hydrogen slightly short ofelectrons, and therefore, making it

slightly positive, just like the picture to the left.Because of this electrical distortion, water is

described as a polar molecule.

When something such as sodium chloride is being dissolved in water, the slightly positive hydrogens cluster aroundthe chlorine, whereas, the slightly negative oxygen cluster around the sodium. The water molecules then literally pullthe sodium chloride crystal apart.

This pull doesnt work on every molecule. Magnesium oxide isnt soluble in water because the water molecules arentstrong enough to break the magnesium-oxygen attractions.Whats so special about methylbenzene?


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Well, methylbenzene is not a polar molecule. It is unable to pull the hydrogen and chlorine apart and therefore,hydrochloric acid wont be formed.

Oxygen and Oxides

(2.15) Composition of Air

This is the approximate composition of air. Memorize it.

There are also very small amounts of noble gases in the air.

(2.16)Showing That Air Contains About 1/5 Oxygen

Using Copper

The apparatus originally contains100cm

3of air. This is pushed

backwards and forwards of the heatedcopper, which turns black as copper(II)oxide is formed. This uses up theoxygen. On cooling, around 79cm

3of

gas is left in the syringes 21% hasbeen used up. Therefore, the aircontains 21% of oxygen.

Using the Rusting of Iron

Iron rusts in damp air, using oxygen up as it does so. The experiment shows some damp iron wool in a test tubecontaining air. The tube is inverted in a beaker of water and the level of the water in the tube is marked by a rubberband. The tube is left for a week or so for the iron to use up the oxygen to makeguessiron oxide!

The water level rises in the tube as the oxygen is used up, and the new level can be marked using a second rubberband. You can find the actual volumes of the gases at the end of the experiment by filling the tube with water to eachof the rubber bands in turn, and pouring it into a measuring cylinder. If the original volume was, say, 15cm

3, and the

final volume was 12cm3, then the oxygen used up measures 3cm

3.

The percentage of oxygen in air was 3/15 x 100 = 20%.

Burning Phosphorus

This can be done by putting a bell jar into a beaker filled with water. Phosphorus on anevaporating dish is placed onto the water (the jar has no bottom). It is then touched with ahot metal rod, which starts the reaction between phosphorus and oxygen. Phosphorus usesup the oxygen to form phosphorus oxide, lowering pressure in the jar and therefore, makingwater levels rise in the jar. The water should rise up by 20%.

Gas Amount in Air (%)

Nitrogen 78.1Oxygen 21.0

Argon 0.9

Carbon Dioxide 0.04


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Equation:

(2.17)Making Oxygen in the Lab

Oxygen is most easily made in the lab from hydrogen peroxide solution using manganese(IV) oxide as a catalyst. Thereaction is known as the catalytic decomposition (splitting up using a catalyst) of hydrogen peroxide.

2H2O2(aq) 2H2O(l) + O2(g)

Reaction of Oxygen with Magnesium, Carbon and Sulfur

Magnesium Magnesium reacts with oxygen to produced white,powdery magnesium oxide. It produces a bright whiteflame during the reaction. It is a base.

2Mg(s) + O2(g) 2MgO(s)

With Sulfur Sulfur burns in oxygen with a tiny blue flame.Poisonous, colourless sulfur dioxide is produced. It isan acidic oxide.

S(s) + O2(g) SO2(g)

With Carbon Carbon burns in oxygen if heated strongly to givecolourless carbon dioxide. Depending on the purity ofthe carbon, a small yellow-orange flame may beproduced.

C(s) + O2(g) CO2(g)

Carbon Dioxide

Preparing It in the Lab

Carbon dioxide is made by the reaction between dilute hydrochloricacid and calcium carbonate in the form of marble chips.

CaCO3(s) + 2HCl(aq) CaCl2(aq) + CO2(g) + H2O(l)

Formation of Carbon Dioxide from Thermal Decomposition ofMetal Carbonates

Key thing here: When heating metal carbonate, you get:

Metal Carbonate Metal Oxide + Carbon Dioxide


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Here is the picture of the experiment setup:

Properties of Carbon Dioxide

Colourless gas, denser than air, slightly soluble in water

Used in carbonated (fizzy) drinks because it dissolves in water under pressure. When bottle is opened,

pressure falls and gas bubbles out of the solution. Used in fire extinguishers to put out electrical fires, or those caused by burning liquids, where using water

could cause problems. The carbon dioxide sinks onto the flames and prevents any more oxygen fromreaching them.

Turns limewater cloudy white (limewater is calcium hydroxide work out the equation yourself water isone of the products).

Carbon Dioxide and Sulfur Dioxide Their Reactions With Water

Carbon Dioxide Carbonic acid is produced when carbondioxide reacts with water. It is a weakacid. This reaction can be reversed bysimply heating or boiling the acid.

CO2(aq) + H2O(l) H+(aq) + HCO3

(aq)

Sulphur Dioxide Sulfur dioxide reacts with water to form

a weak acid known as sulfurous acid,

H2O(l) + SO2(g) H2SO3(aq)

Sulfur Dioxide, Nitrogen Oxide and the Environment

Acid rain is caused when oxygen and water in the atmosphere react with sulfur dioxide to produce sulfuric acid(ouch), or with various oxides of nitrogen to give nitric acid. These mainly come from power stations, burning fossilfuels, motor vehicles etc.

Acid rain can kill trees and make lakes so acidic it cannot supportlife. Limestone and some metals such as iron are also attacked byacid rain.

The solution to acid rain involves removing sulfur from fuels, usingcatalytic converters in cars and scrubbing the gases from power

stations to remove the oxides. The catalyst helps convert nitrogenoxides into harmless nitrogen gas but has no effect on sulfur dioxide.

Methods of Separation

Filtration: For separating an insoluble solid from a liquid, ora soluble solid from an insoluble one.

Sand can be separated from water by pouring the mixturedown a funnel with filter paper. The sand will collect at thefilter paper.

It can also be used to separate sand from something like saltby dissolving salt in water (which leaves you with sand mixedwith salt water). The mixture can then be poured down afunnel. The sand that collects at the top should can be rinsedand dried. The water can be evaporated from the salt byheating with a Bunsen burner. Back

Chromatography: For separating liquids bydissolving them in a solvent. The dyes that make upthe ink differ in two important ways:

How strongly they stick to the paper
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How soluble they are with the solvent

An example would be separating ink colours or plant dyes. A dot of the ink/dye would be drawn onto a piece of paper.It would then be left in water, which acts as a solvent. Because different colours have different solubility levels, somecolours would travel up further on the paper.

Crystallization: Mainly used for purifying substances by forming crystals from a precipitating solution.Crystallization refers to the forming of solid crystals from a homogenous (solution) mixture.

An example would be forming pure salt crystals. This is done by dissolving the impure salt into a solvent such aswater. The salt solution is then allowed to cool. As it does, pure salt crystals would form at the bottom of the water,whereas, the impure substances would be left in the water. The crystals can then be rinsed with a chilled solution anddried.

Distillation: Distillation is good from separatinga liquid from a solution.

An example would be separating water from a saltsolution. The solution would be heated at theliquids boiling point, in this case 100

0C, so it will

leave the solution as a vapour. The vapour would

then condense into a liquid with the help of thecooling water. The vapour, now as a liquid, wouldfall into the beaker.

Fractional Distillation: Fractional Distillationis used to separate two liquids based on theirboiling points.

An example would be separating ethanol from water.Ethanol has a lower boiling water than water (at about78

0C), therefore, the heating is monitored (using the

thermometer) to ensure that the temperature does notreach 100

0C (the boiling point of water). Anyways, the

ethanol would turn into a vapour and travel out of theflask. It would then condense into its liquid form with

the help of the cooling water and fall into the beaker.

Unit 2: Structure and Bonding

Ionic Bonding

Ionic bonding is the bonding in which there has been a transfer of electrons from one atom to another to produceions. The substance is held together by strong electrostatic attractions between positive and negative ions. Ions areformed when it gains or loses electrons. Ones that gain forms negative ions, and ones that lose form positive ions.

A positive ion is called a cation. A negative ion is called an anion.

You can find the charge of an ion by looking at the group it belongs to. If it belongs to groups 1-4, it has a charge of 1-4

+(they are positive), whereas, if it belongs to groups 5-0, it has a charge of 3

0. Below is a table containing charges

of common ionic compounds and transition metals.

Ion Symbol Charge

Silver Ag 1+

Copper (I) Cu 1+


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This is an example of a dot and cross diagram. The crosses represent theelectrons on the sodium (anion) and the dots represent the electrons on thechlorine (cation). In a dot and cross diagram, you must use arrows to showwhich electrons are moved from the anion to the cation. On the finaldiagram, you mark the new electron(s) on the cation as a cross.

Boiling and Melting Points of Ionic Compounds

Ionic compounds have high boiling and melting points due to strong

intermolecular forces between the atoms. This is because when the ionsare formed during an ionic reaction, one of them would be positive, and onewould be negative. Positive and negative attract and therefore, you getsomething like a strong magnet.

As ionic charge increases, so does the melting/boiling points. Ionswith 2

+and 2

-would have stronger attraction because their charges a

stronger, whereas, ions with 1+

and 1-would still have a strong attraction,

but less stronger than 2+-

compounds.

Structure of Ionic Compounds

An ionic crystal consists of giant three-dimensional lattices held together by strong electrostatic attractions betweenthe positive and negative ions.

Structure of Sodium Chloride

This is the basic structure of a sodium chloride crystal. The green is thechloride and the blue is the sodium. Remember that each sodium istouched by six chlorides and each chloride is touched by six sodiums.Look at the middle atoms if unclear. Remember, this structure repeatsitself over and over.

Ionic bonds always produce giant structures.Ions form closely packed regular lattice arrangement.They have high melting/boiling points.The crystals tend to be brittle.

Compounds tend to be soluble in water and insoluble in organicsolvents.

Covalent Bonding

Covalent bonding is formed by sharing a pair of electrons between twoatoms. This is so that both atoms can achieve a full outer shell. It is a strong attraction between the bonding pair ofelectrons and the nuclei of the atoms involved. Covalent compounds are only formed when the reactants are non-metals.

Diagrams YOU Need to Know

Ammonium NH4 1+

Copper (II) Cu 2+

Cobalt Co 2+

Nickel Ni 2+

Zinc Zn 2+

Iron (II) Fe 2

+

Chromium Cr 3

+

Iron (III) Fe 3+

Ion Symbol Charge

Hydroxide OH 1-

Nitrate NO3 1-

Hydrogen Carbonate HCO3 1-

Carbonate CO3 2-

Sulphate SO4 2-

Phosphide P 3-

Phosphate PO4 3-

Nitride N 3-


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Element Diagram Eleme

ntDiagram Element Diagram

H2 CH4 CO2

Cl2 NH3 Ethane

HCl(g) O2

H2O N2 Ethene

Simple Molecular Structures

These are gases, liquids or solids with low melting points. Examples includewater, chlorine, oxygenetcThe covalent bonds between the atoms in a molecule are strong.However, the forces of attraction between these molecules (inter-molecularforces) are weak.They have low melting points, since not a lot of heat is needed to provide the

energy for the molecules to move away from each other, hence, overcome the intermolecular forcesbetween them.They tend to be insoluble in water.They are often soluble in organic solvents.They do not conduct electricity because the molecules have no overall charge and there are no electronsmobile enough to move from molecule to molecule.

Giant Covalent Structures

There are no charged ions.ALL the atoms are joined up to their adjacent atom by extremely strong covalent bonds and packed intogiant regular lattices.They have very high melting points, since a lot of heat is needed to provide the energy to break apart themany strong covalent bonds.They tend to be insoluble in water.They do not conduct electricity.

Diamond

The diamond is the hardest natural substance. It is a form of pure carbon. Each carbonatom forms four covalent bonds to the other carbon atoms. They are arranged in atetrahedral arrangement. Diamond has a very high melting point, obviously due to verystrong carbon-carbon bonds. It does not conduct electricity because all the electrons inthe outer levels of the carbon atoms are tightly bonded between the atoms. None of themare free to move around. Diamond is insoluble like, to both water and other solvents.


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Use of Diamond

Saw blades can be tipped with diamonds in high-speed cutting tools used on stone and concrete. The strongtetrahedral structure makes the diamond hard, making it suitable for this purpose.

Graphite

Graphite is arranged differently it has a layer structure. Each graphite layer is strong,but it is easy to separate individual graphite layers. Each carbon atom only forms threecovalent bonds. Graphite conducts electricity because the fourth electron is free tomove around.

Use of Graphite

Because of the layered structure, graphite can be used as a dry lubricant tolubricate locks.

Metallic Crystals

Metals are giant structures which consist of a regular array of positive ionsin a sea of delocalized electrons. When metal atoms bond together to form

solid, visible metal, their outer electrons are no longer attached to particularelectrons and are free to move around the whole structure.

Metals are able to conduct electricity because the delocalized electrons arefree to move throughout the structure. The energy is picked up by theelectrons and moved around the metals, transferring the electricitythroughout the whole structure. The same goes to heat energy.

Metals are easy to shape because their regular packing makes it simple foratoms to slide over each other. Metals are said to be malleable.

Introduction to Electrolysis

In metals and carbon, electricity and electric current is simply a flow of electrons or ions. Electrolysis is the chemicalchange caused by passing an electric current through a compound which is either molten or in a solution. An

electrolyte is a substance that undergoes electrolysis. It contains ions. It is the movement of the ions, which areresponsible for both the conduction of electricity and the chemical changes that take place. Covalent compounds arenot electrolytes and dont conduct electricity because they have no free moving electrons. Ionic compounds onlyconduct electricity when molten or in a solution because the ions separate and are free to move. These particles canthen carry the electric current.

Diffusion

Experiment to distinguish between electrolytes and non-electrolytes

i. Dissolve substance in water, or if possible, melt it.ii. Put a conductivity tester into the substance.iii. If the light bulb lights up, it is an electrolyte.

Explain: When dissolved in water, free moving electrons are able to carry the electric current across from the

cathode to the anode, completing the circuit and lighting the bulb. If the light bulb does not light up, the substanceis obviously not an electrolyte.

But sugar dissolves, why does the bulb not light up? Sugar is a covalent structure.


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Diffusion happens when particles spread from higher to lower concentration. It requires a concentration gradient).

Potassium Manganate (VII) Experiment

Diffusion through liquids is very slow if the liquid is totally still. This can be shown but dropping a piece of potassiummanganate (VII) into water. It can take days for the colour to spreadbecause the gap between each particle is small.

The Bromine Experiment

Showing diffusion in gases can be done by filling a lower gas jar withbromine gas and topping it with a gas jar filled with air. The bromineparticles and air particles will eventually bounce around to give an evenmixture.

The Ammonium Chloride Experiment

This experiment is used to show that particles in different gases travel atdifferent speeds. It relies on the reaction between ammonia andhydrogen chloride gases to give white solid ammonium chloride.

A white ring of ammonium chloride would form near the hydrochloric acid. This shows that ammonia particles havetravelled further to reach the hydrogen chloride gas, showing that it travels faster.

Dilution

Dilution is the reduction of concentration in a solution.

Showing Dilution and Leading to the Idea of Small SizedParticles

Suppose you dissolve 0.1g of potassium manganate (VII)in 10cm

3of water to give a deep purple solution. Assume

the smallest drop you can see is 1/1000cm3. The whole

solution will be made up of 10000 drops, each dropcontaining 0.00001g of potassium manganate (VII).

Suppose you dilute this down 10 times by taking 1cm3

ofthe solution and making it up to 10cm

3with more water.

Continue doing this until the colour is too faint to see. By

the time of the fifth dilution, each drop will only contain abillionth of a gram of potassium manganate (VII). If you only needed one particle of potassium manganate (VII) perdrop in order to see the colour, the particle cant weigh more than a billionth of a gram.

IS this a good answer? Nowhere near it! A potassium manganate (VII) particle actually weighs about0.00000000000000000000026g! In reality, you need huge numbers of particles in each drop in order to see thecolour.

Dont worry I dont get this either

Unit 3: Organic Chemistry


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Organic chemistry is mainly based around hydrocarbons compounds made only up of hydrogen and carbon. It isdrawn with lines joining carbons and hydrogen. All carbon bonds have to be bonded to hydrogen if not somethingelse. The left picture below shows carbons with all bonds taken up (ethane). The right picture below shows anincorrect picture of a hydrocarbon because one of the carbons has a free bond.

Hydrocarbon compounds that contain carbon and hydrogen only.

Homologous series family of compounds with similar propertiesbecause they have similar bonding. They show a graduation in physicalproperties (mpt/bpt) and similar chemical properties such as the generalformula. Alkanes are the simplest.

Saturated when carbon cannot take anymore bonds singlecarbon-carbon bonds.

Unsaturated presence of a carbon-carbon double bond. General formula The formula of different homologous series of carbons. Isomers molecules with the same molecular formula but different structural formulae.

Learning the Code

Do you have to remember the formula for propane, butane,ethane? No! You can work it out yourself! The first part of thename tells you how many carbons there are in the longest chain(not necessarily in total). By the way you have to learn these atleast the first five. It helps.

For example: propane (left) has three carbons. Butane (right) hasfour carbons.

Alkanes andAlkenes

Alkanes and Alkenes are two homologous series.

Pentene has a five carbon chain with a double bond.

Isomers

We know what isomers are.

Coding for Double Bonds

For things like pentene and butene, there are many places you can put the double bondsin.

Code Number of Carbons

Meth 1

Eth 2

Prop 3

But 4

Pent 5

Hex 6

Hept 7

Oct 8

Non 9

Dec 10

Ending Meaning?

ane All carbon bonds arefilled with hydrogeni.e. they aresaturatedhydrocarbons.

ene There is a doublecarbon-carbon bond

they areunsaturatedhydrocarbons.


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Pent-1-ene means pentene with the double bond on the first carbon-carbon bond (right).Pent-2-ene means pentene with the double bond on the second carbon-carbon bond and so on

But wait! What about pent-4-ene and pent-5-ene? Those dont exist. Why? Because pent-4-ene is pent-2-ene flippedover, and pent-5-ene is pent-1-ene flipped over!

Flip: C-C-C=C-C and you get C-C=C-C-C!

Methyl and Ethyl Groups

If the hydrocarbon has a methyl or ethyl group, these two come first,before the coding for the number of carbons in the chain. But beforeeven the methyl or the ethyl there is a number and hyphen to showwhich carbon has the methyl or ethyl branch.

For example, this is 2-methylbutane

As you can see, there is a methyl group branching off the second carbon. The rulesare similar to double bonds though, there is no such thing as 3-methylbutanebecause that is basically 2-methylbutane flipped over.

But wait! There are five carbons! Why isnt it 2-methylpentane? Because remember,

these names are based on the longest carbon chain in the hydrocarbon and thelongest carbon chain there is 4, hence, butane. This means that 2-methylbutane isan isomer of pentane C5H12.

Some Isomers of Butane C4H10

Some Isomers of Pentane C5H12

Alkanes

Alkanes are a homologous series of saturated hydrocarbons. The first five are methane, ethane, propane, butane andpentane.

Code Meaning

Methyl Has a branch of CH3coming off one of thebonds.

Ethyl Has a branch ofCH3CH2 coming off.


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The general formula for alkanes is: CnH2n + 2

For example, ethane:

Ethane has two carbons, so n=2.The formula of ethane must be C2H2(2) + 2 = C2H6.

Complete Combustion of Alkanes

If there is enough oxygen, alkanes will burn in oxygen completely to give carbon dioxide and water. The generalequation for combustion:

Hydrocarbon + Oxygen Carbon dioxide + Water

The combustion of methane would be: CH4(g) + 202(g) C02(g) + 2H20(l)

Note: Balancing combustion equations can be annoying. An easy way would be to balance them in the order ofcarbon, hydrogen then oxygen.

Incomplete Combustion

If there isnt enough oxygen, you get incomplete combustion, in which carbon monoxide and water are producedinstead. Carbon monoxide is a colourless, odourless and poisonous gas. It is dangerous because it can combine toour haemoglobin and stop it from carrying oxygen. As a result, you get ill or even die because oxygen cannot travel toall parts of your body.

Reaction with Bromine

Alkanes react with bromine under the presence of ultra-violet light. One hydrogen from the hydrocarbon would bereplaced by a bromine atom. This is known as a substitution reaction. Bromine can be used as an indicator foralkanes and alkenes without UV light. Adding bromine water to alkanes produces no colour change. Reactingbromine water to alkenes make it turn from brown to colourless.

However, if the mixture of bromine and methane is reacted under UV light, it loses its colour, a mixture ofbromomethane and hydrogen bromide gases is formed.

CH4(g) + Br2(g) CH3Br (g) + HBr(g)

Alkenes

Alkenes have double bonds, making them unsaturated hydrocarbons.

Alkenes have the general formula of CnH2n the first four being ethene, propene, butene and pentene.

Combustion

Like alkanes, alkenes burn in oxygen or air to give carbon dioxide and water.

Reaction with Bromine

Alkenes undergo addition reactions, in which part of the double bond breaksand is used to join other atoms onto the two carbon atoms. When added toalkenes, and the test tube is shook, the brown of the bromine would bedecolourised, making it suitable as a test for alkenes.


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The product of reacting ethene to bromine gives 1, 2-dibromoethane and is a colourless liquid.

CH2=CH2(g) + Br2(aq) CH2BrCH2Br(l)

Ethanol

All alcohols contain anOH group attached to a carbon chain. Ethanol is C2H5OH.

Production of Ethanol

Hydration of Ethene

Ethanol can be made by reacting ethene with steam (because it contains more energy) a process known as hydration.

CH2=CH2(g) + H2O(g) CH3CH2OH(g)

Only a small portion of ethene reacts. The ethanol is condensed as a liquid and the unreacted ethene is recycled.Explaining the Choice of Temperature

Reversible reactions happen in two ways while ethene is being converted into ethanol, ethanol is also beingconverted back into ethene. Reversible reactions can also shift the equilibriumor alter the reaction. Since thereaction is exothermic the reaction produces lots of heat. If you increase the temperature, the reaction wont like itbecause it is already producing heat, therefore, it would adapt to the conditions by making more ethene so less heatwill be produced. On the contrary, if you decrease the temperature, the reaction would adapt to this by increasingback the temperature; by producing more ethanol in other words, push the equilibrium to the favourable/forwardreaction. However, making the temperature too low would mean super slow reaction, although more ethanol would beproduced. 300 degrees is therefore, a compromise temperature producing an acceptable yield of ethanol in a shorttime.

Explaining High Pressure

In the equation, you have two moles (one mole of ethene and one mole of water) on the left, and one mole (ofethanol) on the right. Increasing the pressure would mean the equilibrium would be shifted forwards. Why? Thereaction would adapt to the conditions by producing more ethanol because you only get one mole of ethanol whichtakes less space than two moles of ethene and water.

Also, theres the collision theory. Increasing the pressure means that thered be less space for the atoms to move.The atoms would also move with more force. This increases the frequency of collisions.

The problem: its expensive and ethene might polymerise and turn into polyethene.

Fermentation

Yeast is added to a sugar or starch solution at 300C for several days in the absence of air for anaerobic respiration.

Enzymes in the yeast lower the activation energy, increasing the rate of conversion of the sugar into ethanol andcarbon dioxide. However, they first have to break the sugars into smaller sugars like glucose. In fact, ethanoic acid isproduced and then converted into ethanol.

For example, sucrose:

C12H22O11(aq) + H20 C6H12O6(aq) + C6H12O6(aq) sucrose + water glucose + fructose

C6H12O6(aq) 2C2H5OH (aq) + 2CO2(g) glucose/fructose ethanol + carbon dioxide

The yeast then gets killed in the mixture, which means that the ethanol produced is impure. To purify it, the alcoholmust undergo fractional distillation.

What is Needed:

Ethene and steam 300

0C

60-70 atmospheres Phosphoric acid

catalyst


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Comparing the two methods

Fermentation Hydration

Use of Resources Uses renewable resources sugar beet orsugar cane, corn and other starchymaterials.

Uses non-renewable resources once oilgets used up, theyre screwed.

Type of Process A batch process everything is mixed andleft for several days. It is then removed anda new reaction is set up quite inefficient.

A continuous flow process a stream ofreactants is constantly passed over thecatalyst more efficient.

Rate of Reaction Slow, takes several days. RapidQuality of Product Produces impure ethanol that needs further

processing.Produces much purer ethanol.

Reaction Conditions Uses gentle temperatures and ordinarypressure relying on anaerobic respirationof yeast.

Uses high temperatures and pressures,needing a high input of energy expensive.

Common Question: Which method would poorer places like Brazil use and why? [3 marks]Answer: Fermentation, because Brazil has the weather conditions to grow large yields of sugar cane and they donthave access to crude oil.

Dehydration of Ethanol into Ethene

Dehydration of ethanol produces ethene and water, using hot aluminium oxide as a catalyst.

CH3CH2OH(g) CH2=CH2(g) + H2O(l)

Crude Oil

Crude oil is a mixture of hydrocarbons. These chains can be super long or super short.

The Trend in Boiling Point and Viscosity

Viscosity means how runny something is Volatile means how easy it turns into vapour at room temperature

As the number of carbon atoms in molecules increases and gets bigger, intermolecular attractions also increase,making it more difficult to pull one molecule away from neighbouring ones. As they get bigger, these changes occur:

Boiling point increases the larger the molecule, the higher the boiling point due to stronger intermolecularattractions.

Liquids become less volatile the bigger the hydrocarbon, the more slowly it evaporates in roomtemperature. This is again, due to strong intermolecular attractions.

Liquids become more viscous (flow less easily) Small hydrocarbons are runny, but large ones are muchstickier and gooey (and viscous) because of intermolecular attractions.

Bigger hydrocarbons do not burn as easily, meaning they are less useful.


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The Fractionating Column

Crude oil is separated in fractionatingcolumn. This process is fractionaldistillation, and splits crude oil into variousfractions depending on their boiling points

and size.

Note: forget about Naphtha

Fraction Uses

Refinerygases

A mixture ofmethane, ethane,propane andbutane.Commonly used fordomestic heatingand cooking.

Gasoline CarsKerosene Used as fuel for jet

aircraft.

As domesticheating oil.As paraffin forsmall heaters andlamps.

Diesel oil For buses, lorries,some cars andrailway engines.

Some is cracked toproduce morepetrol.

Fuel oil For ships Industrial heating

Bitumen Residue from thebottom which can

be used for roads.

Combustion and Incomplete Combustion

Combustion of hydrocarbons produces carbon dioxide and water exothermic. Incomplete combustion of hydrocarbons produces carbon monoxide and water in which carbon monoxide

is dangerous because it can bind to haemoglobin and prevent it from carrying oxygen.


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In car engines, the temperature reached is high enough to allow nitrogen and oxygen from the air to react, formingnitrogen oxides. This contributes to smog and causes irritation to human mucus membranes. As well as that, nitrogenoxides can react with water in the atmosphere and from nitric acid or acid rain.

Cracking

The Crude Oil Problem

Amounts of each fraction you get depend on the proportions of various hydrocarbons in the original crude oil. Farmore petrol is needed, than something like bitumen. In other words, fractional distillation of crude oil produces morelong-chain hydrocarbons than can be used directly, and fewer short-chain hydrocarbons than required.

The solution? Cracking! Cracking is a useful process in which large hydrocarbon molecules are broken into smallerones. Most of the hydrocarbons found in crude oil are long-chain alkanes. Cracking can convert these into alkenesand shorter alkanes. It is an example of thermal decomposition.

How it Works

The fraction is heated to give a gas and is passed over a catalyst of silica or alumina witha temperature of 600-700

oC.

Long alkane alkene + alkane

Sometimes you may get more than one type of alkene/alkane. Make sure the numbers of carbon and hydrogen are balanced.

In an equation, this would read: hexane butane + ethene C6H14 C4H10 + C2H4

Polymers

Alkenes can be used to make polymers. Polymers are big long molecules of single units called monomers. Moleculescontaining carbon-carbon double bonds can be joined together. Part of the double bond is broken and used to join toother monomers. Joining up lots of monomers to make a polymer is called addition polymerisation.

How to Draw a Polymer

Its like drawing a hydrocarbon, except the ends are left blank (so it can join more).

As for the repeating unit (which is the unit that keeps repeating itself), you show thealkene (or the monomer) with its double bond opened up. You then enclose it with brackets andput an n to its right.

Polymers to Know

What is Needed:

Alumina/Silica as

catalyst 600-700

0C


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Polymer Repeating Unit How it looks together Uses

Polyethene Plastic bagsPlastic bottles

Polypropene RopesCrates

Polychloroethene PVC for drainpipesor windowsElectrical insulation

Nylon Condensation Polymer

In condensation polymerisation, when two monomers combine, a small molecule such as water or hydrogen chlorideis lost. Nylon is made through condensation polymerisation.

The two monomers that make up nylon:

Hexanedioic acid

From a family of compounds called dicarboxylic acids.

1,6-Diaminohexane

From a family known as diamines.

Joining them togetherThe lost of a Water Molecule
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As a block diagram (where the (CH2)6 and (CH2)4 become blocks to make it look easier)

End Note: Sometimes you may begiven ClOCCH2 CH2 CH2 CH2COCl instead of hexanedioic acid. In this case, just do the same thing, with the lost of

hydrogen chloride HCl.

Unit 4: Analytical Chemistry and Kinestics

Tests for Ions and Gases

Flame Tests: Taking a piece of nichrome, make loop at the end and dip into salts containing ions


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Ion Colour

Li+

Crimson red

Na+

Yellow orange

K+

Lilac

Ca+

Brick red/orange red

Using Sodium Hydroxide solution

Ion Colour of Precipitate

Cu+

Blue

Fe+

Sludgy Green (or just green)

Fe+

Orange Brown (rust)

For Ammonium Ions (NH4+)

Heat gently and add sodium hydroxide solution. It will give off a distinctive smell of ammonia (NH3). Ammonia can betested by holding a damp red litmus paper. Since it is alkaline, it will turn damp red litmus paper from red to blue.

Using Dilute Nitric Acid and Silver Nitrate Solution

Ion Colour of Precipitate

Cl-

White

Br-

Pale Cream

I-

Yellow

For Sulphate Ions (SO42-

)

Using dilute hydrochloric acid solution and then adding barium chloride solution to form a white precipitate of bariumsulphate.

For Carbonate Ions (CO32-

)

Using dilute hydrochloric acid to react with the carbonate, to produce carbon dioxide gas which can be tested bybubbling through limewater, turning it from colourless to cloudy, milky white.

Tests for Gases

Gas Test Result

Hydrogen Hold a lit splint in presence of hydrogen gas. Produces a squeaky pop.

Oxygen Hold a glowing splint in presence of oxygen gas. Glowing splint relights.

Carbon Dioxide Bubble through limewater. Turns limewater from colourless to cloudy,milky white.

Ammonia Hold damp red litmus paper in ammonia gas. Turns damp red litmus paper blue.

Chlorine Hold damp blue litmus paper in chlorine gas. Bleaches or turns blue litmus paper white.

Solubility Patterns

All nitrates are soluble. All sodium, potassium and ammonium compounds are soluble.


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Most carbonates and hydroxides are insoluble except for sodium, potassium and ammonium. All sulphates are soluble except barium and lead(II) sulphate. All chlorides are soluble except lead(II) and silver chloride.

Reactions of Metals to Acids

Metals react very similarly to dilute hydrochloric acid and dilute sulphuric acid.

Metals Reaction to Acid

Magnesium Rapid fizzing, mixture gets very hot, colourless magnesium sulphate/chloride solution forms.

Aluminium Is slow due to its coat of aluminium oxide which prevents aluminium from contacting the acid.On heating, this layer is removed, aluminium will start fizzing rapidly abit like Mg.

Zinc Zinc reacts slowly with cold dilute acid and may produce some effervescence. On heatinghowever, it fizzes more.

Iron Iron also reacts slowly with cold dilute acid and will produce abit of effervescence when heated.

Combustion of Hydrogen

Bonds are broken in the hydrogen and oxygen molecules. These form new bonds of water molecules. This reaction isexothermic, and gives out water in the form of steam, before it condenses into a liquid. The reaction is:

2H2(g) + O2(g) 2H2O(l)

Testing for Water

Water turns white anhydrous copper(II) sulphate blue. Its reaction is

CuSO4(s) + 5H2O(l) CuSO45H2O

Or you can use cobalt chloride paper which turns pink in the presence of water.

You can check the purity of water by showing that it freezes at exactly 0C and boils at exactly 100C.

Rates of Reactions

Experiment Setup

To measure the effects of changes in surface area, concentration of solutions, temperature and use of catalyst, youcan react calcium carbonate marble chips with dilute hydrochloric acid and measure the mass of CO2 produced byweighing the difference in mass of the reactants and the mass of the products (there wont be any change in massproduced, because the initial mass of reaction will equal the final mass, however, since carbon dioxide gas is formed,this will escape from the flask, and therefore, the amount of mass lost will bethe mass of carbon dioxide produced. Plot the results on a graph with massagainst time and youll get an upward curve.

For the reaction to occur, acid particles must collide with the surface of themarble chips. As the acid particles get used up, the collision rate decreases, sothe reaction slows down.

Changes in Surface Area of Solid

You can repeat the above experiment by keeping thesame mass of marble chips, just using smaller ones toincrease the surface area. The reaction happensfaster. You have to remember the graphs. Noticehowever, that in the end, the amount of carbon dioxideproduced is still the same just that the small chipsexperiment happens faster.


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Why does it happen faster? Because the surface area in contact with the gas or liquid is much greater. Less marblechip particles are hidden away from the acid particles.

Changes in the Concentration of Solutions

Repeat the original experiment but using hydrochloric acid

only half as concentrated as before. The graph should looksomething like this (ignore the 80% line) in which thereaction happens slower and produces half as muchcarbon dioxide gas:

In terms of collision theory, if you increase theconcentration of reactants, the reaction becomes fasterbecause it increases the frequency of collisions persecond.

Changes in the Temperature of the Reaction

Do the original experiment again, but this time, at a higher temperature.Your graph will look like this (ignore the concentration label cause itsWRONG unless the lower concentration solution is still in excess).

Increasing the temperature means more kinetic energy for the particles,which make them move faster, therefore, making them collide morefrequently.

Also, not all collisions make new bonds. Some particles just bounce off eachother. In order for a reaction to happen, particles have to collide with aminimum amount of energy called activation energy. Increasing thetemperature produces a very large increase in the number of collisions thathave enough energy for a reaction to occur.

In the following diagram, a) shows a fail collision and b) shows a successfulone.

Changingthe Pressure

Changing the pressure of a reaction where the reactants are onlysolids or liquids makes virtually no difference, so the graphs remainunchanged. Increasing the pressure in a reaction where the reactantsare gases does speed the reaction up. This is because it forces theparticles closer together, so they hit each other more frequently.

Catalysts and HowThey Work

Catalysts speed up the rate of reactions but arent used up in theprocess. You can show that manganese (VI) oxide is a catalyst bysimply having two conical fasks containing hydrogen peroxide.


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Hydrogen peroxide decomposes to give oxygen and water. Put the manganese (VI) oxide in one of the flasks.Oxygen would be given off quickly. To check that the manganese (VI) oxide hasnt been used up, simply filter it outfrom the solution and weigh it (remember to weigh it before the experiment too!). The graph should look like thepressure graph in which the rate of reaction increases, but the amount you get at the end is still the same.

So how does it work?

Adding a catalyst gives the reaction an alternative route for reactions with a lower activation energy.

Unit 5: Quantitative Chemistry and Energetics

A mole is a measure of the amount of substance. One mole contains 6 x 1023

(also known as the Avogadro Number)particles (atoms, molecules or formulae) of the substance. For example, 1 mol of sodium contains 6 x 10

23atoms of

sodium.

Calculating Relative Atomic Mass

Chlorine has two isotopes: chlorine-35 and chlorine-37. A typical sample will be 75% chlorine-35 and 25% chlorine-37.

The RAM = (0.75 x 35) + (0.25 x 37) = 35.5g

Calculating Relative Formula Mass

Straightforward stuff.

Amount in moles = Mass of Substance (g) /RFM of Element or Compound (g)

Calculations using moles:

The equation for sodium chloride is: 2Na + Cl2 2NaCl

If 2.3g of Na was used:

a) Find out how many moles of Chlorine was usedb) Find out the volume of Chlorine used in the reactionc) Find out the mass of sodium chloride produced

a) Firstly, convert the grams of Na into moles:2.3 / 23g = 0.1 mol

The equation says that 2 moles of Na and 1 mole of Cl (1 mole of a diatomic molecule is always X2) is needed toproduce 2 moles of NaCl, so if 0.1 mol of Na is used, then half of that is the amount of chlorine used in the reaction inmoles.

So moles of Cl used = 0.1 / 2 = 0.05 mol

b) One mole of any gas has a volume of 24 dm3

(24000cm3) at room temperature and pressure. This is also

called the molar volume.

Cl2 is a gas and the moles used in the reaction = 0.05 molSo the volume of Cl2 gas used = 0.05 x 24000 = 1200cm

3

c) The moles of NaCl produced is 0.1 mol (if 2 moles of Na gives 2 moles of NaCl, then 0.1 mole of Na will give0.1 mole of NaCl). So all you do is:i) Find the RFM of NaCl (58.5)ii) Multiply that by 0.1 (5.85g)

Molar Concentrations The Hard Part

H2 H x 2 2 x 1 = 2

Ca(OH)2 (1 x Ca) + (2 x O) + (2 x H) 40 + (2 x 16) + (2 x 1) = 74


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Remember that:

Mol/dm3

means moles per litre (e.g. a salt solution of 0.5 mol/dm3

means 0.5 moles (or 58.5/2 = 29.25g) ofsalt was dissolved in a litre of water

Its all about proportion

20 cm3

of 0.5 mol dm3sodium hydroxide solution was dissolved with 25cm

3of hydrochloric acid to form a sodium

chloride solution. Calculate the concentration of HCl needed to react with the NaOH

NaOH + HCl H2O + NaCl

RFM of NaOH = 40g = 1 mole of NaOH0.5 mol dm

3of NaOH means (40 x 0.5) 20g of NaOH was dissolved in 1000cm

3of water

The amount of moles in 20cm3

of NaOH solution:

20cm3/1000cm

3x 0.5 moles = 1/50 x 0.5 = 0.01 moles of NaOH

The equation says that 1 mole of NaOH + 1 mole of HCl gives 1 mole of NaClSo 0.01 moles of NaOH + 0.01 moles of HCl gives 0.01 moles of NaCl

So 0.01 moles of HCl was present in 25cm3

of HCl solution! However, concentration is measured in mol dm3

so:

1000cm3/25cm

3x 0.01 mol = 0.4 mol dm

3of HCl used.

Calculating the Empirical Formula and Molecular Formula

The empirical formula is the simplest formula and only tells you theratio of the various atoms. Suppose 2.4g of magnesium combinedwith 1.6g of oxygen, you can use a table to work out the empiricalformula. (Mg = 24 O = 16)

What about withpercentagefigures?

Suppose you had a compound containing 85.7% C, 14.3% H andyou were asked to calculate the empirical formula. Firstly, youassume that 100% = 100g! (C = 12 H = 1)

However, you know that CH2 does not exist. Remember this is only the ratio. To find the molecular formulae, youneed to know the relative formula mass of the compound. Suppose it was 56g for the above question.

Firstly, find out the RFM of CH2 = 12 + 2 = 14gFind out how many times 14 goes into 56, so 56/14 = 4 timesWhich means the molecular formula is C4H8!

Obtaining Formulae Experimentally

Metal Oxides

Hydrogen can be passed over metal oxides to reduce it to themetal. To find the formula of copper oxide, the experimental stepsare as follows:

1. Measure the mass of the empty combustion tube.2. Use a spatula to put copper oxide into the tube. Weigh

the tube.3. Set up the apparatus as shown. Turn the gas at the jet

to light the excess gas.

Mg O

Combining Masses 2.4 1.6

Number of moles 2.4/24 1.6/16

= 0.10 0.10

Ratio of Moles 1:1

Empirical Formula MgOC H

Percentage 87.5 14.3



= 7.14 14.3

Ratio of Moles 1:2Empirical Formula CH2


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4. Heat the copper oxide until it has all turned into red copper.5. Stop heating but leave gas passing through until everything has cooled.6. Weigh the combustion tube.7. Put masses in a table and calculate empirical formula from there.

In the Case of Water of Crystallisation

When substances crystallise from a solution, water becomes chemically bounded with the salt. This is called water ofcrystallisation and the salt is said to be hydrated.

Suppose you had to find the formula of a BaCl2nH2O (a barium chloride crystal), to find n:

1. Weight the mass of an empty crucible.2. Add barium chloride crystals and reweigh.3. Heat the crucible gently (so the barium chloride wont decompose), and reweigh.

4. Put masses into a table and calculate the formula from there.

Calculating Percentage Yield

Most of the time, when you do carry out a chemical reaction, you get less than you expect. The rest of it has been lostin some way perhaps due to spillages or losses when chemicals are transferred.

Suppose you work out that 10g of A will give 500g of the product, but you only get 400g?

The percentage yield is (400/500) x 100 = 80%

A general formula would be: (mass produced/expected mass to be produced) x 100

Endothermic and Exothermic Reactions

H represents the molar enthalpy change forexothermic and endothermic reactions

Mass of Empty Tube 52.2g

Mass of tube + Copper Oxide (Before) 66.6g

Mass of tube + Copper 65.0Mass of Oxygen 66.6 65.0 = 1.6g

Mass of Copper 65.0 52.2 = 12.8g

Cu O



= 0.20 0.10

Ratio of Moles 2:1

Empirical Formula Cu2O

Mass of Empty Crucible 30.00g

Mass of tube + Crystals (Before) 32.44g

Mass of tube + Anhydrous Crystals(After)

32.08g

Mass of BaCl2 2.08g

Mass of Water 0.36g

BaCl2 H2O



= 0.01 0.02

Ratio of Moles 1:2

Empirical Formula BaCl22H2O

Endothermic Heat energy istaken inBreaking of bonds

H = + N kJ mol-

Exothermic Heat energy isgiven outMaking of bonds

H = - N kJ mol

-


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Energy Calculations

The general formula:

Bonds of all the reactants Bonds of all the products = Energy change

Example: Methane reacts with chlorine to produce chloromethane and hydrogen chloride. The equation:

CH4 + Cl2 CH3Cl + HCl

You would be given a table with the bonds and the energyrequired to break/bond them:

Reactants:

4 C H bonds (CH4) = 4 x 413 = 1652kJ1 Cl Cl bond (Cl2) = 1 x 243 = 243 kJ Total: 1652 + 243 = 1895 kJ

Products:

3 C H bonds = 3 x 413 = 1236 kJ

1 C Cl bond = 1 x 346 = 346 kJ1 H Cl (HCl) = 1 x 432 = 432 kJ Total: 2017 kJ

(Carbon can form 4 bonds. In this case, 3 of them bonds with 3 hydrogen and the last one bonds with chlorine)

Energy Change = 1895 2017 = -122 kJ the reaction is exothermic

Describing Simple Calorimetry Experiments

All these involve measuring a temperature change during the reaction. Specific heat is the amount of heat needed to raise the temperature of 1g of a substance by 1

0C. For water,

the value is 4.18 J g-1

0C

-1(joules per gram per degree Celsius).

Heat Given Out = Mass x Specific Heat x Temperature Rise

For Neutralisation, Displacement Reactions and Dissolving

Bond C - H C - Cl H - Cl Cl - Cl

Energy (kJ mol-

) 413 346 432 243

Mass of Weighing Bottle + Mg (g) 10.810

Mass of Weighing Bottle Afterwards (empty) (g) 10.687

Mass of Mg used (g) 0.123


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They all follow the same method. This exampleinvolves measuring the heat evolved (or energy)when magnesium reacts with dilute sulphuric acid.

1. Pour an excess of sulphuric acid into a polystyrene cup and measure the temperature of the acid.2. Pour some magnesium powder into a weighing bottle and weight it.3. Pour the powder into the acid and record the highest temperature.

4. Weigh the empty weighing bottle.

Lets say the total mass of the solution and Mg is 50g

Heat evolved when 0.123g of Mg reacts = 50 x 4.18 x 10.1J = 2111J = 2.111kJ

To find out the heat evolved when 1 mole of Mg reacts (Mg = 24g):

(2.111/0.123g) x 24 = 412 kJ

The temperature rose, meaning the reaction is exothermic so:

Mg(s) + H2SO4(aq) MgSO4(aq) + H2(g) H = -412 kJ mol-1

This is actually smaller than the accepted value, which is around -417 kJ mol-1

. A reason for this could be that heat

was lost too quickly. Using a mercury thermometer may give better results.

Combustion

1. Put 100cm3

of water into a conical flask and record thetemperature.

2. Fill the spirit burner with alcohol (lets say ethanol) and weight.3. Light the spirit burner and record the temperature of water until

there is say, a 400C increase.

4. Reweigh the spirit burner.

Volume of water (cm ) 100

Mass of water being heated (g) 100

Mass of burner before (g) 37.355

Mass of burner after (g) 36.575

Mass of ethanol burnt (g) 0.780

Original temp. of water ( C) 21.5

Final temp. of water ( C) 62.8

Water temperature increase ( C) 41.3

Heat gained = 100 x 4.18 x 41.3 = 17260 J = 17.26 kJ

Ethanol is C2H5OHOne mole of ethanol = 46g

Amount of heat produced from 1 mole of ethanol = (17.26/0.780) x 46 = 1020 kJ

KEY POINTS FOR THIS UNIT

1 mole is the Avogadro constant number of particles The molar volume is 24 dm

3or 24 000 cm

3

Number of moles = mass / RFM 1 dm

3= 1000 cm

3

Mol dm-3

= mol per 1000 cm3

Initial Temp. ( C) 17.4

Final Temp. ( C) 27.5

Temperature Rise ( C) 10.1


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G dm-3

= grams per 1000 cm3

Energy/Heat = mass x specific heat x temperature rise

Unit 6: Chemistry in Society

Electrolysis

RECAP! Electrolysis is a chemical change caused by passing an electric current through a compound which is eithermolten or in solution. An electric current (in chemistry terms), is a flow of electrons or ions. An electrolyte is a

substance that undergoes electrolysis. Electrolytes all contain ions. Ionic compounds, for example, are electrolytes.Electrolytes can only undergo electrolysis when molten or in a solution, where the ions are free to move. Covalentcompounds are not electrolytes because they dont contain ions.

Electrolysis can form new substances when ionic compounds conductelectricity. It is set up as so:

The electrodes are usually made of carbon because it isfairly un-reactive.

The positive electrode is called the anode. The negative electrode is called the cathode.

A simple example of electrolysis involves molten (melted) lead (II)bromide:

So what happens?

Molten lead is found at the bottom of thecathode.

Bromine gas comes out of the anode.

When the power supply is switched off, no morebubbles are produced and everything elsestops.

What the hell happened?

Since the lead (II) bromide (PbBr2) is molten, itsions are free to move around.

The bromide ions are attracted to the positiveelectrode. The extra electron which makes thebromide ion negatively charged is deposited intothe anode, thus, turning them back into neutral


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bromine atoms. These then covalently bond to join bromine atoms (i.e. bromine gas).

On the other hand, the lead ions gain back to electrons (it has a 2+

charge) and become normal lead atoms.These fall to the bottom of the container as molten lead.

The half-equation at the anode would be:

2Br

-

Br2 + 2e

-

What it basically means is that, two bromine ions are formed whenone bromine molecule receives two electrons (to fill its shell). This isthe format for all anions (ions that are negatively charged).

The half equation atthe cathode wouldbe:

Pb2+

+ 2e- Pb

All cation half-equations are of this form. Half-equations basicallyshow the gaining and losing of electrons. This basically means thatthe lead (II) ions get two electrons to become a neutral lead atom.

With molten substances, the metal will be produced at the cathodeand whatever its bonded to will be produced at the anode.

When you electrolyse aqueous solutions (not molten salts), things are much different because you have to considerthe water molecules too. Water is a weak electrolyte but it can ionise to form hydrogen and hydroxide ions.

If the metal is more reactive than hydrogen, thenhydrogen ions from water is discharged instead. Thesepair up to form hydrogen gas that escapes as bubbles.

If the metal is below hydrogen, you get the metal

produced.

If you have solutions of halides (chlorides, bromides oriodides), you get the halogen (chlorine, bromine oriodine) produced.

With other negative ions such as sulphates, oxygenwould be produced.

The electrolysis of sodium chloride solution (brine) does not give sodium and chlorine! Heres the electrolysis:


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Sodium is higher than hydrogen in the reactivity series, so

hydrogen ions from the water in the sodium chloride solutionis discharged instead at the cathode.

2H+

+ 2e- H2

Chloride ions give up one electron each (chloride ion = 1-charge) and become chlorine atoms. These covalently bondto form chlorine gas and bubbles out of the solution at theanode.

2Cl- Cl2 + 2e

-

When all the chlorine has been removed from the solution,only hydroxide (OH

-) ions and sodium (Na

+) ions are left, as

well as some water. These combine to form sodiumhydroxide solution (NaOH).

The electrolysis of sodium chloride solution is used to manufacture sodium hydroxide solution. The process is slightlydifferent it is electrolysed in a diaphragm cell:

The products are kept separated by the diaphragm. If thechlorine produced were to react to hydrogen, it would cause anexplosion on exposure to sunlight or heat to give hydrogenchloride. Furthermore, if the chlorine were to react with thesodium hydroxide solution formed, it would form bleach. Uses ofsodium hydroxide include:

Making bleach Making soap Making paper NaOH breaks the wood down

Uses of chlorine include:

Sterilising water Making hydrochloric acid Making bleach

And the electrolysis of copper sulphate solution:

Copper is lowerthan hydrogen and

therefore, a coat of it forms at the cathode.

Cu2+

+ 2e- Cu

Oxygen gas is discharged from the hydroxide ions in the waterbecause the sulphate ions are more stable.

4OH- 2H2O + O2 + 4e

-

If you electrolyse the solution for longer, something else happens. The

hydrogen ions are being discharged and remains in the solution. Similarly,sulphate ions are being discharged either. As a result, the solution turns into sulphuric acid (H2SO4) and it beginselectrolysing:

Sulphate ions are being discharged from the acidso oxygen is discharged from the hydroxide ionsinstead.

4OH- 2H2O + O2 + 4e

-


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There are only hydrogen ions arriving at the cathode so they discharge as hydrogen gas.

2H++ 2e

- H2

Common Exam Question: Why is twice as much hydrogen produced than oxygen? For every four electrons that flow

around the circuit, one molecule of oxygen and two molecules of hydrogen are produced.

Electrolysis Calculations

Back to moles! Here are some things to know:

One faraday means one mole of electrons passing around the circuit. One faraday = 96000 coulombs. Charge (coulombs) = Current (amps) x Time (seconds)

Example: What mass of copper is deposited on the cathode during the electrolysis of copper (II) sulphate solution if0.15A flows for 10mins?

The electrode equation is:

Cu2+ + 2e- Cu

Calculate the coulombs involved:

10mins x 60 = 600 secondsCharge = 0.15 x 600 = 90 coulombs

The equation says that 1 mole of copper ions + 2 moles of electrons give 1 mole of copper atoms

1 mole of electrons = 96000 coulombs2 moles of electrons = 192 000 coulombs2 moles of electrons (192 000 coulombs) give 1 mole of copper (RFM = 64g), so 90 coulombs give:

(90/192 000) x 64g = 0.03g

(Coulombs worked out/coulombs of electrons) x RFM of element = Mass of element deposited

When involving gases

Example: During the electrolysis of dilute sulphuric acid, hydrogen is released at the cathode and oxygen at theanode. Calculate the volumes of hydrogen and oxygen produced if 1.0A flows for 20mins

The electrode equations are:

2H++ 2e

- H2

4OH- 2H2O + O2 + 4e

-

Assume the molar volume of gas to be 24000 cm3

For hydrogen:

2H++ 2e

- H2

2 hydrogen ions + 2 moles of electrons give 1 mole of hydrogen molecule

20mins x 60 = 1200 seconds1200 x 1A = 1200 coulombs

2 x 96 000 coulombs (2 moles of electrons) = 192 000 coulombs1 mole of hydrogen gas = 24000 cm

3


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192 000 coulombs give 1 mole of hydrogen gas (or 24000 cm3

of hydrogen gas)So 1200 coulombs give: (1200/192 000) x 24000 = 150 cm

3of hydrogen produced

For oxygen:

4OH- 2H2O + O2 + 4e

-

20mins x 60 = 1200 seconds1200 x 1A = 1200 coulombs

4 moles of hydroxide ions give 2 moles of water + 1 mole of oxygen gas + 4 moles of electrons (4 x 96000 = 384 000coulombs)

1 mole of oxygen gas = 24000 cm3

384 000 coulombs give 1 mole of oxygen gas (or 24000 cm3

of oxygen gas)So 1200 coulombs give: (1200/384 000) x 24000 = 75 cm

3of oxygen produced

The equation:

(Calculated coulombs/moles of electrons) x (mole of gas x 24000) = amount of gas produced in cm3

Reversible Reactions and Dynamic Equilibria

Some reactions are reversible. Reversible reactions are indicated by the symbol

Some examples of reversible reactions include:

Copper (II) Sulphate Crystals

Heating the blue hydrated copper (II) sulphate crystalscauses them to lose their water of crystallisation, makingthem turn from blue to white the white copper (II)sulphate crystals are described as anhydrous meaningwithout water:

CuSO45H2O CuSO4 + 5H2O

However, this reaction can be reversed by simply adding water to thecrystals. The crystals will become hydrated again:

CuSO4 + 5H2O CuSO45H2O

Heating Ammonium Chloride

When ammonium chloride is heated in a test tube, the white crystals decompose into

hydrogen chloride gas and ammonia gas. These flow upwards and recombine againfurther up the test tube:

NH4Cl HCl + NH3

This later recombines:

HCl + NH3 NH4Cl

Introducing Dynamic Equilibria


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Things change when reversible reactions are carried out under closed conditions meaning no substances areadded to the reaction mixture and no substances can escape from it. Heat however, can be given off or absorbed.

In a reversible reaction, you have the forward reaction (the reaction going from left to right) and the back reaction (theopposite of the forward reaction) happening at the same time. Both rates of reactions will become equal and this pointis the dynamic equilibrium. It is dynamic in a sense that the reactions are still continuing, and equilibrium in a sense

that the total amounts of the various things present are now constant. In other words:

A + 2B C + D

When you have a reaction like the above, A + 2B (forward reaction) is reacting to produce C + D (back reaction). Atthe same time, C + D is reacting to produce A +2B. In the end, you have equal amounts of products and reactants.

Another way to think of is, is to imagining walking down an elevator that goes up, making sure youre walking at thesame speed as the elevator. You would be going down, but everytime you take one step down, the elevator goes onestep up. In the end, you remain where you are.So how would you produce more of substance C in a reversible reaction such as the above? You can do this byaltering the position of the equilibrium by either:

Changing the pressure Changing the temperature Increasing/Decreasing the concentrations of substances present

Adding a catalyst

If a dynamic equilibrium is disturbed by changing the conditions, the reaction moves to counteract the change.

In other words, the reaction will either go more towards the forward direction or the back direction in an attempt toadapt to the conditions.

A + 2B C + D

Changing the Concentration

What happens when more of A is added? If you add more A, the reaction will wan

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The Ultimate IGCSE Guide to Chemistry