HINCHINGBROOKE SCHOOL

CHEMISTRY DEPARTMENT A BRIDGING COURSE TO SIXTH FORM CHEMISTRY

Inspiring Excellence Fulfilling Potential

INTRODUCTION A-LEVEL CHEMISTRY

Before you start studying A-level Chemistry at Hinchingbrooke you should have completed this new bridging course. It has been specifically written to help you prepare for your AS studies. It explains the principles from GCSE and then extends this to help you start thinking like an A-level Chemist. Over the summer holidays, you will need to find time to work through this booklet. It is vital that you become confident in the concepts covered as they are the principles on which the A-level course is based. Here at Hinchingbrooke we are following the AQA course. It is important that you are willing and committed to completing this course. One of the most important skills you will have to develop for all A-levels is independent learning. The content of this book will be assessed over the first half term. It is important to be willing to put the effort into practicing and applying the concepts to the exercises enclosed. With commitment and practice of the concepts that are encountered in Chemistry, this will lead you to the essential ingredient for success – confidence in the ability to learn. As you work through this book there are exercises to be completed, websites to try out and videos to watch. The answers to the exercises are enclosed at the back of the book. However, cheating won’t help your understanding! A-level Chemistry is a difficult and complex subject. If you are committed to learn and pass at AS level, you will acquire useful knowledge and fantastic transferable skills. Good luck in working through this booklet and we will see you in September.

ATOMIC STRUCTURE A CLOSER LOOK AT ATO MS

An atom is believed to be made up of three sub-atomic particles, the proton, the neutron and the electron. The properties of which can be summarised in the table below: Question 1 Complete the table Particle

Relative mass

Relative charge

Location

proton

neutron

electron

The locations of these sub atomic particles are arranged as shown below in a Lithium atom:

The atom is fundamental part of all elements. It is no longer the smallest particle that can exist but it is the smallest particle that still retains the properties of the element. AS indicated about the electrons stay in shells around the nucleus. In principle Chemistry is the study of what happens to electrons in an atom. At Alevel Chemistry they are the focus of our studies.

THE PERIODIC TABLE The Periodic table contains all the information that we need to work out the structure of an atom. It will allow us to predict the nature of an element and so understand its properties. The periodic table is available to you in all the written exams. It is therefore fundamental that you know what information it contains and how to use it. The diagram below has the periodic table shaded different colours.

Question 2 Indicate in the box below what each of the colours represents: Colour

Represents

Within each box of an element in the periodic table are three pieces of information. 1. The elements symbol, e.g. Al for Aluminium. 2. The smallest number – the atomic number, which is the number of protons within an elements nucleus e.g. 13. As elements have no overall charge, the number of electrons must be the same as the protons. 3. The biggest number the mass number. This is the total number of protons and neutrons in a nucleus e.g. 27.

All the information is located in the Periodic table. Using the GCSE Period table below complete the following table:

Question 3

ELECTRONIC CONFIGURA TION

The arrangement of the electrons around the nucleus at GCSE uses the Bohr model. Within this model each of the shells can only take a certain number of electrons. In order of increasing distance from the nucleus, these are: 

First shell – up to 2



Second shell – up to 8



Third shell - up to 8

It is key to remember that the first shell is always filled before the second and so on.

Question 4 Complete the diagrams of the following elements:

Carbon

Oxygen

Magnesium

Name the following atoms:

ISOTOPES When looking at the Periodic table there are two versions. On page 5 the atomic mass is rounded to the nearest whole number, with the exception of chlorine and copper. The table on page 4 has all the numbers between 1 and 2 decimal places. If we look specifically at chlorine, it has an atomic mass of 35.5. As a result it is difficult to calculate the number of neutrons, as you cannot get half a neutron. Instead this is showing the relative atomic mass. This takes into account that there are a number of different atoms of an element which have a different number of neutrons. In chlorines case, there are still 17 protons in each atom type but in 25% of the atoms on the planet chlorine has 20 neutrons with an atomic mass of 37. In the other 75% there are 18 neutrons with a total mass of 35. This means that the total relative atomic mass of four chlorines atoms picked at random would be: (3 x 35) + (1 x 37) = 105 + 37 = 142 142 / 4 = 35.5 Atoms of the same element which have a different number of neutrons are called isotopes. Question 5 Neon has isotopes and relative atomic mass?

. They exist in the ≈ ratio of 9:1 respectively. What would be the average

BONDING Elements will react with other elements to form compounds. The way in which they react is linked to their electrons. At GCSE we explain that the elements react in one of two ways, as you will see it is not that straight forward. It is true to say they combine in one of two ways and this is the next fundamental idea A-level chemistry is based on. The two ways are ionic and covalent bonding.

IONIC BONDING Atoms are stable when they have a full outer shell. The properties of an element with a full outer shell are that they are unreactive. Neon is unreactive as it has a full outer shell. Let’s look at the reaction of sodium and chlorine. Sodium is a metal in Group 1 (the alkali metals). It has one electron in its out shell. If it gives this electron away it will have a full outer shell, which means it is stable. If sodium has lost an electron it also means it now has a charge. There are now more protons in the nucleus than electrons around the outside. AS a result the atom has now got a positive charge. It can no longer be called an atom, it is an ion. Chlorine on the other hand is a non-metal in Group 7 (the halogens). It has seven electrons in its outer shell. If it gain one it to will have a full outer shell and become stable. In chlorine’s case, it gains an electron and so has more electrons than protons. This results in an overall charge of minus 1. As it has a negative charge it is now a negative ion. Ionic bonding involves the transfer of electrons from a metal ion (sodium for example) to a non-metal (e.g. chlorine). When the electron is transferred the metal atom becomes charged and is called a positive ion. The non-metal becomes charged and is a negative ion. Both sets of ions now have a full outer shell of electrons, just like the noble gas atoms. A diagram showing the reaction of sodium and chlorine is below:

It is vital to fully understand the reactions, the movement of electrons and the resulting ion formation. In the boxes below draw similar diagrams to NaCl on the previous page for the stated compounds: Question 6 Reaction of to form Lithium Fluoride

Reaction to form Magnesium oxide

PROPERTIES OF COMPOU NDS WITH IONIC BONDI NG It is key to remember principle properties of ionic compounds: 1. They have giant ionic structure 2. They have high melting points and boiling points 3. Ionic solids cannot conduct electricity 4. If ionic solids are molten or dissolved they can conduct electricity 5. Many ionic compounds are soluble in water Questions 7 1. Why do ionic compounds have high melting points? 2. Why do ionic solids not conduct electricity?

COVALENT BONDING Non-metals cannot give away electrons to gain full outer shells. Instead they must have an alternative method to resolve the lack of electrons if they are to react with each other. Let’s focus on Hydrogen. We know from GCSE that hydrogen is a diatomic molecule that exists in the air as H2. This means that a hydrogen atom has reacted with itself to form a more stable molecule than a single hydrogen atom. The diagram showing the position of the electrons is:

Each atom has donated an electron and the two atoms are now sharing the electrons. One atom donated the red electron the other the blue. In comparison to the ionic bonding, in covalent bonding: 

There is no transfer of electrons



There are no charged particles



The covalent bond is between only the atoms sharing the pair of electrons



There is no strong attraction between molecules (intermolecular)

In the boxes below draw similar diagrams to H2 on the previous page for the stated compounds: Question 8 Reaction of to form Water

Reaction to form Carbon dioxide

PROPERTIES OF SIMPLE COVALENT MOLECULES

1. Most are gases or liquids at room temperature 2. Those that are solid have quite low melting points 3. Simple molecules do not conduct electricity 4. The molecules do not have any overall charge. 5. Most often the molecules do not dissolve in water; if they do they do not conduct electricity.

PROPERTIES O F GIANT COVALENT STRUCTURES There are some specific examples that are needed which are giant covalent structures. At GCSE these are; 

Silicon dioxide



Diamond



Graphite

DIAMOND Diamond is an allotrope of the element carbon. It is a giant covalent structure. Each carbon is attached to four other carbon atoms by strong covalent bonds. Diamond has a very high melting point and is very hard. The structure of Diamond

GRAPHITE Graphite is also an allotrope of carbon. Each carbon is bonded to another 3 carbon atoms in the same layer by strong covalent bonds. Graphite has a high melting point. Its layers can easily slip over each other ans so it is slippery and soft. Graphite can conduct electricity. The structure of graphite

Question 9 1. Why can graphite conduct electricity? 2. What is an allotrope? 3. What is the Mohs scale, and why is diamond significant on it? 4. In the box below draw the structure of silicon dioxide and explain its key properties

DRAWING COVALENT MOLECULES AS indicated in the structures of diamond and graphite. It is not possible to always draw the overlapping shells in covalent compounds. Taking the original example of H2 we can now revisit how covalent compounds are drawn. The bonding seen before in H2 was:

This is more commonly draw as:

H–H The single line represents a single covalent bond. H2 is a linear molecule. This means that when drawn the two hydrogens are drawn in the same plane. It is not always the case that the sticks are a straight line like in H2 . Water for example has a bent shape, as you will see in the AS lessons this is all due to the electrons on oxygen. Its structure is:

At GCSE there were a set of covalent molecules whose structure was required as recall. This remains the case at A-level. In the boxes below draw the stick structures of the following compounds: Question 10 Methane, CH4

Hydrogen chloride, HCl

Ammonia, NH3

Carbon dioxide, CO2

FORMULAE A FUNDAMENTAL SECTIO N YOU MUST UNDERSTAND Question 11 Before you go any further make sure you can define the following terms: a) Atom b) Element c) Compound d) Ion You must know the following section really well, this will allow you to access all of the concepts in the Alevel chemistry course. The formula of a compound is useful as it indicates how many particles are present in one entity of a compound. An entity is the simplest formula unit of a compound, e.g. H 2O, NaCl or Li2CO3. It is at this point we must look back at the Periodic table, and focus on the groups:

Group

1 +1 H Li Na K

2 +2 Positive ions Be Mg Ca

3 +3

B Al

4 + or -4

C Si

5 -3

N P

6 7 -2 -1 Negative ions O S

F Cl

0 0 He Ne Ar

When looking at constructing the formula of a compound it is important to look at the group and work out what ion the element will form. When elements react they will gain, lose or share a specific number of electrons until they have a full outer shell. All compounds have no overall charge. With this in mind let’s look at the formula of some compounds. Sodium chloride We showed the formation of this ionic compound on page 7. It was clear that sodium had to get rid of one electron and chlorine had to gain one. As a result both gained a full outer shell. As the exchange of electrons was 1:1 the formula of the compound is then the symbols in a 1:1 ration i.e. NaCl. It is also critical to note that sodium formed a Na+ ion and Chlorine a Cl- ion. When the two charges are added together the overall charge is 0. Magnesium oxide This ionic compound is composed of Magnesium and oxygen. Magnesium is in group two and forms a 2+ ion. Oxygen is in group 6 and as a result forms a negative 2- ion. Like in NaCl, these two react in a 1:1 ratio. This means the formula for magnesium oxide is MgO.

Sodium oxide In this case the 1+ ion of sodium reacts with the 2- ion of oxygen. To make these have an overall charge of 0, we need two sodium ions. As a result the reaction is in a 2:1, metal to non-metal ratio. This means the formula for Sodium oxide is Na2O. Question 12 Complete the table below: Reacting elements

Formula of compound

Name of compound

Potassium and Fluorine Calcium and sulphur Lithium and nitrogen Aluminium and oxygen

It is also important to remember that there are some transition metals that at GCSE you needed to know what their most common ion is. The table below indicates the metals and their charges, some old ones, some new ones: +1 Ag+ Au+ Cu+

+2 Positive ions Cu2+ Zn2+ Co2+ Mn2+ Fe2+

+3 Fe3+ Cr3+

WHEN TO USE BRACKETS At times there is a need to use brackets. A small subscript number after the bracket multiplies all the elements in the bracket by that number. The brackets are needed when groups of elements go together and form an ion. In GCSE there are specific ones to remember, these are still needed for A-level. 1- ions

2- ions Ion

3- ions

Ion NO3-

Name Nitrate

OH-

Hydroxide

CO32-

Carbonate

MnO4-

Manganate

CrO42-

Chromate

Cr2O72-

Dichromate

SO4

2-

Name Sulfate

Ion PO

3-

1+ ions

Name Phosphate

Ion NH4+

Name Ammonium

Magnesium hydroxide In this compound the 2+ ion of magnesium reacts with the 1- ion of hydroxide. To balance out the charges there must be two OH- ions to react with one Mg2+. As a result of this the formula of magnesium hydroxide is Mg(OH)2. It cannot be MgOH2 as this would indicate there are one Mg, one O and two H. This is not the case and so the brackets are essential to identify the compound correctly. Calcium carbonate There are times when the brackets are not needed and this is when the grouped element ions are in a 1:1 ratio. The 2+ ion of Calcium reacts with the 2- ion of carbonate. This is in a one to one ratio. AS a result the formula is CaCO3, with no requirement for the brackets. Question 13

Reacting element/group Sodium and nitrate Aluminium and hydroxide Potassium and managanate Ammonium and sulfate

Formula of compound

Name of compound

EQUATIONS WORD AND SYMBOL EQUATIONS In chemistry the word equation tells us what chemicals will be reacting and the products that they form. The balanced symbol equations have a completely different function. They tell us what in proportion the substances react together. Question 14 What do the following symbols represent in an equation? (s) (l) (g) (aq) A balanced symbol equation has the same number of each type of atom on either side of the arrow. Chemical reactions just simply rearrange the atoms; nothing is lost so both sides must have the same number of each atom within it. If balancing equations remains a concern for you, do not worry. We will go over this in September. In the meantime here is a link to a video on You Tube that some of our students have found useful in explaining how to do it: http://www.youtube.com/watch?v=_B735turDoM Here are a summary of hints to remember when balancing equations 1. You must not change the formulas of the compounds/elements involved 2. You can only change the number of atoms by putting numbers in front of the formulae. 3. You will quite often need to go back and forth between the sides until you get it right. 4. If an equation is impossible to balance – go back and check the formulas this is often the cause. Question 15 Balance the following equations 1) HCl + NaOH  NaCl + H2O

6)

Cr + HCl  CrCl3 + H2

2) SO2 + O2  SO3

7)

Fe3O4 + H2  Fe + H2O

3) Fe2O3 + C  Fe + CO

8)

C3H8 + O2  CO2 + H2O

4) Fe2O3 + CO  Fe + CO2

9)

H2 + CuO  Cu + H2O

5) NH3 + O2  NO + H2O

CHEMICAL CALCULATION S RELATIVE MOLECULAR MASS One of the most universally used equations in chemistry is that which calculates the relative mass of a given compound. As you have probably already noticed, terminology in Chemistry is vital. You will be asked to work out the relative molecular mass or the relative formula mass of a compound. The key thing to remember is, it is asking the same thing – molecular mass is of a covalent compound, the formula mass is of an ionic compound. The need for two different terms we will discuss in September Either way the equation is exactly the same: Relative molecular mass = sum of the relative atomic mass of all the elements in the compound Mr = Σ Ar As always there are parts along the way which may catch you out: Brackets Make sure you multiply all the elements in the brackets by the little number outside to obtain the total mass e.g. Mg(OH)2 Relative formula mass (Mr) = Mg + (2x O) + (2 x H) = 12 + 32 + 2 = 46 Water of Crystallisation These compounds will become more common in A-level chemistry. They are the ones which have the .xH2O at the end of the formula. These are hydrated salts and have water associated with them. The most common example you have met to date is hydrated copper sulfate CuSO 4.5H2O. This is the blue crystal you have seen since Year 7. There is an anhydrous form of CuSO4 and this is a very pale whitish blue powder, nothing like the bright blue crystals. When calculating the relative formula mass of the compounds with water of crystallisation you must remember to add the correct number of water molecules to the compound e.g.

Molecular mass of H2O = (1 x 2) + 16 = 18 Formula mass of CuSO4 = 63.5 + 32 + (4 x 16) = 159.5 In hydrated copper sulfate there are 5 molecules of water, therefore Formula mass of CuSO4.5H2O = 159.5 + (5 x 18) = 249.5

Question 16 Calculate the mass of the following formulas: Formula

Mr

Formula

Cu(NO3)2

CoCl2.6H2O

C 2 H6 O

(NH4)2SO4

ZnCl2.4H2O

Al2(SO4)3

CO(NH2)2

CuCO3

Mr

THE MOLE Let’s look at the combustion of methane: CH4 + 2O2  2H2O + CO2 If we calculate the masses of the reactants and products we notice that the two sides equal each other e.g. Mr CH4 = 16 Mr O2 = 32 Mr H2O = 18 Mr CO2 = 44 When we take into consideration the balancing the sum of reactants is:

Σ reactants = 16 + (2 x 32) = 80 Likewise for the products:

Σ products = (2x 18) + 44 = 80 This example shows us that if we consider the mass of the reactants and products they react in rather complicated proportions e.g. 16:64:36:44. Chemists needed an easier way to compare quantities and this is where the mole becomes important. Just like a dozen is 12, a score is 20 and a trio is 3. A mole is a specific number – Avogadro’s number to be precise: A mole = 6.022 x 1023 Avogadro worked out how many atoms of carbon-12 were in 12g of carbon. As a result it was 6.022 x 1023. We now have that a mole is 6.022 x 1023 particles of a species (molecule, atom, ions, and electrons) to allow comparison for what takes part in a chemical reaction. In A-level chemistry we will need to use this number to work out reacting masses and concentrations.

CALCULATIONS USING M OLES You have encountered moles at GCSE, the formula for calculating moles is: Moles = mass / Mr When calculating the number of moles of a substance you have it is key to remember that the mass of 1 mole of anything is equal to its molecular mass e.g. 1 mole of water has a mass of 18g, 2 moles of carbon dioxide has a mass of 88g. Question 17 Using the equation above, calculate the number of moles in each case: 1. 2. 3. 4. 5. 6.

640 g of sulphur dioxide 3.2g of sulphur dioxide 0.896 g of sulphur dioxide (to 3 s.f.) 225 g of scandium 0.0045 g of scandium 6.37 g of scandium (to 3 s.f.)

USING THE CORRECT UNITS There will be a requirement to complete a significant number of mathematical calculations throughout you’re a-level course. One of the biggest barriers for any of the calculations is making sure the numbers are in the correct unit for a given equation. It is essential that you know how to convert from one unit to another. Thermodynamic temperature A number of equations in chemistry need to take into consideration the temperature the reaction occurs at. Kelvin proposed a scale that had no negative values. A temperature quoted in Kelvin is the thermodynamic (absolute) temperature. The scale was set putting absolute zero as 0. We however on a daily basis measure in Celsius (oC). Absolute zero is -273.15oC. To convert from oC to K all you need to do is add 273 (the 0.15 is not needed at AS level). This means that -100oC in Kelvin is: -100 + 273 = 173K Question 18 Convert these temperatures to Kelvin: 1. 150oC 2. 2654 oC 3. -59 oC

Units of Volume Volume is measured in the Science lab in cm3. However, this unit is rearly used in equations. It is therefore important to be able to convert from one unit to another. As volume is a cubed unit, it must be converted in 3 directions. So the conversions are: mm3 ÷ 1000 cm3 ÷ 1000 dm3 ÷ 1000 m3 Question 19 Convert the following into dm3 1. 2. 3. 4.

25cm3 150mm3 37cm3 500cm3

Solutions and their concentrations When a substance is dissolved we need to know its concentration to be able to calculate the number of moles involved. Likewise to know the concentration we need to know the number of moles. The equation you met at GCSE is still as vital at A-level: Concentration (moldm-3) = moles ÷ volume (dm3) Question 20 Calculate the following: 1. 10g NaCl dissolved in 100cm3 of solution gives what concentration? 2. 0.16g of KOH dissolved in 250cm3 of solution gives what concentration? 3. How much anhydrous copper sulfate do you need to dissolve in 500cm 3 to obtain a concentration of 0.1moldm-3? 4. How many moles of KMnO4 are there in 500 cm3, if the concentration is 1.5 moldm-3? 5. How much NaOH is needed to dissolve in 350cm3 to obtain a concentration of 0.5 moldm-3?

FINAL THOUGHTS A-LEVEL CHEMISTRY AT HINCIHINGBROOKE You should have worked through this book over the summer break and now have a good foundation on which to start building you A-level knowledge. A-level Chemistry is a difficult subject to study. With perseverance and determination you can master the concepts within the course. As you continue through the course the calculations will become familiar and easier, the concepts more logical and your analytical skills more refined. We will go over all the concepts encountered in this book and ensure that you have fully understood them in September. We will do this before we embark on the AS content. We look forward to meeting you in September. Mrs Alford

ANSWERS 1. Particle

Relative mass

Relative charge

Location

proton

1

+1

Nucleus

neutron

1

0

Nucleus

electron

1/1840 (≈ 0)

-1

Orbiting the nucleus

2. Colour

Represents Gases Metals Transition metals Metalloids

3.

C N

12

7

7

Mg P

12 16

31

Ar Li

6

7

12

12

15 18

18

22

3

3

4

4.

Sodium

5. Relative atomic mass of neon =(9 x 20) + (1 x 22) = 180 + 22 = 202/10 =20.2

6. Reaction of to form Lithium Fluoride

Chlorine

Oxygen

Reaction to form Magnesium oxide

7. a) Ionic compounds have high melting points because a lot of energy must be supplied to overcome the strong forces of attraction between the ions. b) The ions are not free to move and so the solids cannot conduct electricity.

8. Reaction of to form Water

Reaction to form Carbon dioxide

9. a) Why can graphite conduct electricity? It contains free electrons within its layers b) What is an allotrope? Different forms of the same element that exists in the same physical state. c) What is the Mohs scale, and why is diamond significant on it? A scale of hardness used to classify materials. It runs from 1 to 10. Diamond is 10 on the scale and the hardest know natural substance. d) In the box below draw the structure of silicon dioxide and explain its key properties

10. Methane, CH4

Hydrogen chloride, HCl

Ammonia, NH3

Carbon dioxide, CO2

11. a) Atom – smallest particle of element that has the properties of that element. b) Element – a substance that contains atoms which all have the same atomic number c) Compound - a substance formed when two or more chemical elements are chemically bonded together d) Ion – a charged particle formed when an atom or molecule gains or loses one or more electrons

12. Reacting elements

Formula of compound

Name of compound

Potassium and Fluorine

KF

Potassium fluoride

Calcium and sulphur

CaS

Calcium sulphide

Lithium and nitrogen

Li3N

Lithium nitride

Aluminium and oxygen

Al2O3

Aluminium oxide

13. Reacting element/group

Formula of compound

Name of compound

Sodium and nitrate

NaNO3

Sodium nitrate

Aluminium and hydroxide

Al(OH)3

Aluminium hydroxide

Potassium and managanate

KMnO4

Potassium Manganate (VII)

Ammonium and sulfate

(NH4)2SO4

Ammonium sulfate

14. (s) - solid (l) - liquid (g) - gas (aq) – aqueous (dissolved in water)

Question 15 1) HCl + NaOH  NaCl + H2O already balanced 2) 2SO2 + O2  2SO3 3) Fe2O3 + 3C  2Fe + 3CO 4) Fe2O3 + 3CO  2Fe + 3CO2 5) 2NH3 + 2½ O2  2NO + 3H2O multiples allowed 6) 2Cr + 6HCl  2CrCl3 + 3H2 7) Fe3O4 + 4H2  3Fe + 4H2O 8) C3H8 + 5O2  3CO2 + 4H2O 9) H2 + CuO  Cu + H2O already balanced

Question 16 Formula

Mr

Formula

Mr

Cu(NO3)2

187.5

CoCl2.6H2O

237.9

C 2 H6 O

46

(NH4)2SO4

132

ZnCl2.4H2O

208.4

Al2(SO4)3

342

CO(NH2)2

50

CuCO3

123.5

Question 17 1.

640 g of sulphur dioxide = 10 moles

2.

3.2g of sulphur dioxide = 0.05 moles

3.

0.896 g of sulphur dioxide (to 3 s.f.) = 1.40 x 10-2 moles

4.

225 g of scandium= 5 moles

5.

0.0045 g of scandium= 1.00 x 10-4 moles

6.

6.37 g of scandium (to 3 s.f.) = 1.42 x 10-1 moles

Question 18 1. 150oC = 423K 2. 2654 oC = 2927K 3. -59 oC = 214K

Question 19 Convert the following into dm3 1. 2. 3. 4.

25cm3 = 0.025 dm3 150mm3 = 1.50 x 10-4 dm3 37cm3= 0.037 dm3 500cm3= 0.5 dm3

Question 20 Calculate the following: 1. 10g NaCl dissolved in 100cm3 of solution gives what concentration? = 1.71 moldm-3 2. 0.16g of KOH dissolved in 250cm3 of solution gives what concentration? = 1.14 x 10-2 moldm-3 3. How much anhydrous copper sulfate do you need to dissolve in 500cm 3 to obtain a concentration of 0.1moldm-3? = 7.975g 4. How many moles of KMnO4 are there in 500 cm3, if the concentration is 1.5 moldm-3? = 0.75 moles 5. How much NaOH is needed to dissolve in 350cm3 to obtain a concentration of 0.5 moldm-3? = 7g