structure and bonding - studytime.co.nz · the magic that happens when atoms decide sharing is...

NCEA | Walkthrough GuideLevel 2CHEMISTRY

11

Hydrogen

1.00794

H

321

Litium

6.941

Li

11281

Sodium

22.98976928

Na

192881

Potassium

39.0983

K

37281881

281882

28181881

2818321881

2818321882

28181882

Rubidium

85.4678

Rb

55

Caesium

132.9054519

Cs

87

Fransium

(223)

Fr

422

Beryllium

9.012182

Be

12282

Magnesium

24.3050

Mg

202882

Calsium

40.078

Ca

38

Strontium

87.62

Sr

56

Barium

137.327

Ba

88

Radium

(226)

Ra

212892

Scandium

44.955912

Sc

39281892

2818102

281832102

28183232102

2818122

281832112

28183232112

28181892

28181992

2818321892

28183218102

Yttrium

88.90585

Y

2228102

Titanium

47.867

Ti

40

Zirconium

91.224

Zr

72

Hafnium

178.49

Hf

104

Rutherfordium

(261)

Rf

2328112

Vanadium

50.9415

V

41

Niobium

92.90638

Nb

73

Tantalum

180.94788

Ta

105

Dubnium

(262)

Db

57

Lantanum

138.90547

La

89

Actinium

(227)

Ac

58

Cerium

140.116

Ce

90

Thorium

232.03806

Th

CH

NN N

O Ph

STRUCTURE AND BONDING

Introduction 3The Basics 4

Atomic Structure 4Valence Electrons and Stability 5Covalent Bonds 6Intermolecular Forces 7

Lewis Diagrams 8Molecule Shape 11

VSEPR Theory 12The Molecular Shapes 13

Polarity 17

Defining Polarity 17Electronegativity 18Non-polar Covalent Bonds 20Polar Covalent Bonds 21Polarity of Molecules 22

Types of Solids 25

Introduction to Solid Properties 26Molecular Solid Structure 28Molecular Solid Properties 30Ionic Solid Structure 32Ionic Solid Properties 33Metallic Solids 36Metallic Solid Properties 37Covalent Network Solids 38Diamond 38Graphite 39Silicon Dioxide (Silica) 40

Thermochemistry 42

Enthalpy and Enthalpy Change 42Endothermic Reactions 43Exothermic Reactions 46Activation Energy 47Enthalpy Calculations Part 1 48Enthalpy Calculations Part 2 50

Key Terms 53

Level 2 Chemistry | Structure and bonding

Level 2 Chemistry - Structure and bonding | © Inspiration Education Limited 2017. All rights reserved.3


INTRODUCTIONThis standard is all about The Force…

Actually, it’s not the kind of force Jedi use, so put those lightsabers away.

Instead, this standard is all about intermolecular and intramolecular forces. It doesn’t matter if you don’t know what those are just yet, you’ll know soon enough, and you’ll know how these two seemingly simple things make glass different from diamond and plastic bottles radically different from aluminium cans.

Join us on a journey beginning with the humble atom and ending with explaining why metals can conduct electricity while diamond can’t, or why you can dissolve table salt in water but there’s no way you can splash some water on a car and make it disappear.

What will you learn in this walkthrough guide?

Kicking off this cram guide is a “back to the basics” section looking at all the stuff you hopefully didn’t forget from last year. It is essentially looking at the humble atom, what it’s made of, and why it’s important.

Then the real fun begins by looking at molecules and the shapes that they form. From here we look at a very important thing called “polarity”, which has a massive impact on many of the molecule’s properties. After this we move onto molecular, ionic, metallic, and covalent network solids, where you will need to be able to describe the structure and bonding that occurs, as well as the properties that each solid has. The last part of this standard is thermochemistry, which looks at the energy changes that occur in all of these chemical reactions.

So, take a deep breath and dive in!

A word on exam strategy.

This standard can be quite daunting at first as there is A LOT of content to cover! The good thing is that each part of this topic builds on the last, giving you one big picture in the end. The fundamental idea behind this external is the relationship between structure and function - if you understand the structure of molecules or other solids you should be able to figure out their properties, and if you know the properties you can guess what the structure could be.

Here at StudyTime, we’re pretty much GCs (good citizens), so to help you out, we’ve made this guide in plain English as much as we can. We’ve also included a glossary for some of the key terms that you’ll need to master for your exam.

4 Level 2 Chemistry - Structure and bonding | © Inspiration Education Limited 2017. All rights reserved.


However, the language we use isn’t always something you can directly write in yourexam. When this is the case, we offer a more scientific definition or explanation (in ahandy blue box) underneath. These boxes are trickier to understand on your first readthrough, but contain language you are allowed to write in your exam. Look out forthem to make sure you stay on target!

THE BASICSLet’s all be honest; we’re only studying chemistry to blow stuff up and make pretty colours using hopefully dangerous chemicals. All of these awesome reactions are just atoms and molecules breaking and forming chemical bonds. But before we can understand the super duper fun chemistry, we need to take a step back and have a look at:

Those teensy tiny atoms and what they look like when you brutally crack ‘em open. Why we only care about the outermost electrons. The magic that happens when atoms decide sharing is caring. Intermolecular forces: a sciency-sounding word that makes you seem like you know what you’re talking about when you use it.

Atomic Structure

Have a think back to level 1, and try to remember that atoms are made up of:

Positively-charged protons and neutral neutrons inside a densely packed area called the ‘nucleus’.Negatively-charged electrons that zip around the nucleus.

Electronshells

Proton (p+)Neutron (n0)

Electron (e-)

Nucleus

The electrons are more organised than any NCEA student; they sort themselves out into different electron shells, or orbitals. Things get very cosy in the 1st electron shell with just 2 electrons allowed. Once we get to the 2nd and 3rd electron shell, there can be up to 8 electrons.



The outermost electrons often feel a bit neglected, so we give them a special name: valence electrons. This is because the outermost occupied electron shell is called the valence electron shell.

Valence shellNucleus Valence electron

Electrons

-

-

-

-- - - -

--

--

--

- -

- -

-

-

--

-

-

-

- - -

--

--

--

--

STOP AND CHECK:

Turn your book over and see if you can remember:

The structure of an atom. How electrons are arranged around the nucleus. What the valence electron shell is.

Try explain it in your own words.

Valence Electrons and Stability

Atoms may be stable or unstable, and it all depends on their valence electron shell.

Atoms ‘want’ to be stable, and for stability they need a full valence shell

Noble Gases, right at the end of the Periodic Table, have full valence shells.

However, every other atom only has a partially-filled valence shell. Therefore, they have to find some way to get a full valence shell and gain that stability.

Metal atoms (the ones in the first two columns of the Periodic Table, as well as aluminium, this year) are lazy.

They only have 1 or 2 (or 3 for aluminium) valence electrons but it doesn’t make much sense for them to find any more to fill up their valence electron shell, so instead they just toss the ones they do have to a non-metal atom nearby. We can’t blame them though; they’re just taking the easiest route to getting a full valence shell – we’d all do the same!



e

e

e

e

e

e

e

e

e e

e

e

e

e

e

e

e

e

e ee

eIʼm out

valence shellvalenceelectron full valence shell

loses valence electron

+

Non-metal atoms (towards the right of the Periodic Table) are the opposite.

They are so close to having a full valence electron shell and this frustrates them. They hang around and scavenge any electrons donated from metal atoms.

unfilled valence shell full valence shell

gains electron

-

e

e

e

e

e

e

e

e

e

e e

e

e

e

e

e

e

e

e

e e

Remember, atoms are neutral, having the same number of protons and electrons

If an atom gains electrons there will be more negative charges than positive charges, making it a stable, negatively-charged ion, which we call anions. But if an atom loses electrons there will be less negative charges than positive charges, making it a stable, positively-charged ion, which we call cations.

STOP AND CHECK:


Why atoms ‘want’ a full valence electron shell. How cations and anions are formed.

Try to explain it in your own words.

Covalent Bonds

Sometimes a non-metal atom can’t find a metal atom to steal valence electrons from. But as the saying goes: “sharing is caring”. So, they turn to another non-metal atom and



Covalent Bonds

share valence electrons so that both will end up with full valence electron shells.

A covalent bond forms when these two atoms share valence electrons...

When there are at least two atoms connected via these covalent bonds, a molecule forms. These covalent bonds, which hold atoms within individual molecules together, are called intramolecular forces.

It’s a good habit to start thinking of covalent bonds as ‘intramolecular’ forces - especially because their opposites are about to come in very handy.

covalent bonds intramolecular force

hydrogen atom

He

eH

STOP AND CHECK:


Why covalent bonds are formed. How covalent bonds are formed.


Intermolecular Forces

Now that we have talked about intramolecular forces, we need to talk about intermolecular forces. Like the Inter-islander ferry, which travels between the islands of New Zealand, The “inter-” part of the name suggests that these are forces between molecules.

Think about what happens when water freezes into ice. As we are still dealing with water molecules, we can’t gain any new intramolecular forces, however, the closer packing of molecules tells us something else must be going on. In fact, it all has to do with the forces between those water molecules.

Unlike the atoms within each molecule, which are held together by intramolecular forces (covalent bonds), molecules themselves are held together by forces called van der Waals forces. These are examples of intermolecular forces.



We can look at the strength of an intermolecular force by looking at a substance’s melting/boiling point. As we go from solid to liquid to gas, intermolecular forces are the ones which need to be broken to free up these molecules - we will get into this more later.

intermolecular force1 H2 molecule

H He

eH H

e

e

H H

e

eH H

ee

STOP AND CHECK:


What intermolecular forces are.


Quick Questions

What are the different parts of an atom? Why do atoms ‘want’ a full valence shell? How can they obtain a full valence shell? What are covalent bonds and why are they formed? How are molecules held together in the liquid and solid state?

LEWIS DIAGRAMSNow that we’ve sorted out what makes an intramolecular force different from an intermolecular one, let’s work out how to represent them.

Lewis structures, or Lewis diagrams, show the position of atoms within a molecule and the covalent bonds that connect them.

O

?



They look a little something like this:

orFBFF

FBFF

We represent the electrons using either dots (like above) or crosses. Covalent bonds can be shown as two electrons between atoms (like on the left) or the two electrons can be replaced with a line (like on the right).

This may look like a lot to take in, but we promise the following steps should help you to draw any Lewis diagram:

1. First, we need to work out how many valence electrons we have to play with. To do this, count the number of valence electrons in each atom in the molecule. Add these all up to find the total number of valence electrons in the molecule.

2. Next up, we need to work out which of the atoms in the molecule is the “central atom”. This is important as it will be the atom that requires the most bonds. How do we know which one this is? It is the atom that is furthest from having a full valence shell - i.e. a valence electron number furthest away from either 0 or 8.

3. Now that we have our central atom, we need to connect the others to it. We do this by placing a covalent bond, which is represented by two valence electrons, between the central atom and each of the outer atoms.

4. Most atoms obey the Octet Rule where they must have 8 valence electrons in order to be stable. In Level 2, the exceptions to this rule are hydrogen, which only needs 2, beryllium, which only needs 4, and boron, which only needs 6. So, once you’ve added your bonds, you want to add electrons around the central atom and outer atoms so that they all have full valence electrons.

5. Now, you need to check how many electrons you have allocated. You have to use all of the total number of electrons you calculated in Step 1.

6. If you have used more electrons than you started with there’s a bit of a problem. In this case you need to form double bonds between the central atom and some of the outer atoms.

a. If you have used 2 more electrons that what you started with you need to make an additional bond. b. If you have used 4 more electrons that what you started with you need to make two additional bonds.c. The number of additional bonds is equal to the difference divided by 2.

But, running through an example is much more useful.



Consider the molecule ammonia, NH3

Nitrogen, N, is in Group 15 of the Periodic Table and so has 5 valence electrons. Each hydrogen, H, has just 1 valence electron.

The total number of valence electrons in the molecule is: 5 + 1 + 1 + 1 = 8.

Our central atom is going to be nitrogen as it is the furthest from having a full valence shell. We place nitrogen in the centre and place the hydrogen atoms around it. Our next step is to place a single covalent bond between nitrogen and each of the hydrogen atoms:

H XX NH

HXX

XX

Great! Now it’s time to make sure each atom has a full valence shell. The hydrogen atoms only need 2 electrons so they’re already sorted. Nitrogen, however, needs 8. It only has 6 around it at the moment so we just place two more electrons above it.

H XX NH

HXX

XX

XX

So far we have used 8 valence electrons (3 covalent bonds + 1 non-bonding pair) which means we have used them all up. Looks like we’re done - that’s the Lewis structure for ammonia.

Consider the molecule carbon dioxide, CO2

Carbon, C, is in Group 14 of the Periodic Table and so has 4 valence electrons. Each oxygen, O, has 6 valence electrons.

The total number of valence electrons in the molecule is: 4 + 6 + 6 = 16.

Our central atom is going to be carbon as it is the furthest from having a full valence shell. We place carbon in the centre and place the oxygen atoms around it. Our next step is to place a single covalent bond between carbon and each of the oxygen atoms:

O XX C OX

X

Great! Now it’s time to make sure each atom has a full valence shell. All of our atoms need 8 electrons in total. Carbon currently has 4 around it so we place an extra 4,



while oxygens only have 2 so need 6 more.

O XX X

X X

X X

X X

X X

X XX

XX

XXC OX

X

So far we have used 20 valence electrons (2 covalent bonds + 8 non-bonding pairs) which means we have used more electrons than we started with (we only started with 16). If we have used 4 extra electrons we need 2 more covalent bonds. Therefore, let’s try make a double bond between the carbon and each of the oxygen atoms:

O XX

XX

XXC OX

X

Just like before we can make sure each atom has a full valence shell:

O XX

XX X X X

X X X XX

XXC OX

X

Checking again, we have used 16 valence electrons (4 covalent bonds + 4 non-bonding pairs) which means we have used all the electrons up. Looks like we’re done - that’s the Lewis structure for carbon dioxide.

Quick Questions

Use the steps above, as well as a Periodic Table, to draw the Lewis diagrams for:

Water (H2O) Methane (CH4) Oxygen (O2)

MOLECULE SHAPESo, we have seen how all the atoms in covalent molecules bond together, but now we need to think about where they all fit. In Level 2 Chemistry it’ll come out as one of five shapes: linear, bent/v-shaped, trigonal planar, trigonal pyramidal, and tetrahedral.

In this section, we’re going to cover:

VSEPR Theory – what it stands for and what it means for our molecules. How we get each of the molecular shapes mentioned above.

?



VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) theory can be used to predict the shapes of molecules

It sounds quite daunting, but stick with us.

Remember that, as like charges repel, electrons want to move away from one another. As they do this, they end up in a number of different shapes - depending on how many are there to repel each other.

Repel

Attract+ -+ +- -

In level 2, these can be summarised into five different shapes - which can in turn be predicted by VSEPR theory!

VSEPR theory is all about it electrons. In fact, when it comes to molecular shapes they’re the only ones who matter.

The first thing to remember, is that electrons always work in pairs.

These pairs come in two flavours:

1. Bonding pairs (shared between two atoms and make up the covalent bonds) 2. Non-bonding (lone) pairs.

bonding (shared) pairs of electrons

non-bonding (Ione)pairs of electronsx x

x xx x

x x

xx

xx

To determine the shape of a molecule, you need to do some counting

You must count the number of bonding pairs of electrons (remember, these are the covalent bonds, where double and triple covalent bonds can be counted as just 1),



and the number of non-bonding (lone) pairs of electrons around the central atom.

lone pair

lone pair

NH HH

But what about everything else? At this point, you can forget about the outer atoms the central atom is bound to.

Once you have the right combination of bonding and non-bonding pairs of electrons use the table in the next section to get the right shape. It’s a really, really good idea to know these shapes well, otherwise you’ll have major regrets sitting in the exam. In the next section we’ll quickly go over why these shapes are the way they are.

STOP AND CHECK:


How we can use VSEPR Theory to figure out molecular shapes. The two types of electron pairs around an atom in a molecule.


The Molecular Shapes

Ultimately, the shape of the molecule comes comes down to the repulsion of negative charges around the central atom.

This repulsion comes from both bonding pairs of electrons in the covalent bonds between atoms, or the non-bonding (lone) pairs of electrons around the central atom. As “like repels like”, these negative charges try to get as far away from each other possible.

Depending on the combination of bonding and non-bonding pairs of electrons, we get left with different arrangements

The easiest way to picture this, is to imagine a central atom. Now, picture two more atoms joining onto it. If you wanted them to be as far away from each other as possible, where would you put them?



If you answered in a line, you’re correct! Our first shape is linear - which, as the name suggests, just places the central atom and two other atoms in a line.

Because we know from maths that angles on a straight line are 180°, we know the angles between each bond will all be 180°.

180°

O C O

Now, imagine bringing in another group - which we can’t see as a bond, but exists as a lone pair. This new group repels and ‘pushes’ the two bonded pairs of electrons down, resulting in a ‘bent’ or ‘v-shaped’ molecule.

Because they sit closer to the central atom than the bonded atoms, the lone pair actually exerts slightly more repulsion on the two bonded pairs we had earlier. This means that the bond angles in the ‘bent’ shape are just less than 120°.

<120°

1 lone pairrepulsion

OS

O

If we have two lone pairs, like we do in water, the bond angle is even smaller. That’s because the extra lone pair adds more repulsion, pushing the bonding pairs even closer together.

104.5°

more repulsion

2 lone pairs

HO

H

Now, imagine if you swapped the lone pair out for another bonded pair. You would get three bonds evenly distributed around the central atom. This results in the ‘trigonal planar’ shape - with bond angles of 120°.



B

F

F F

120°Boron TrifluorideBF3

Once again, we’ll stir everything up by adding another lone pair to the mix. A fourth electron pair, in the form of a lone pair pushes down on our three bonds - causing them to fold into the ‘trigonal pyramidal’ shape with a bond angle of less than 109.5°.

NH

HH

<109.5°

Last but not least, have a go at replacing the lone pair with one last bond - to form a grand total of four bonding pairs. This gives us the tetrahedral shape, and four evenly distributed bonds at angles of 109.5° from each other.

CH

HH

H109.5°

Because these shapes are three-dimensional, we have a bit of a problem when we try to draw them out on our two-dimensional paper.

Luckily, there are some special conventions which come to the rescue and show how to draw molecular shapes on paper:

A solid line (like on a lewis diagram) represents a bond running along the paperA dashed line represents a bond going into the paperA ‘wedge’ shape represents a bond coming out of the paper - and towards the viewer



These may seem a little odd to begin with, but the more you draw them, the more you will get used to them!

The following table gives the list of molecular shapes that you need to know for Chemistry 2.4:

Bonding Pairs of Electrons

Non-bonding (Lone) Pairs of Electrons

Molecular Shape Name

Bonding Angle

Example

2 0 Linear 180°x x

xx

x x

xx O = C = O

2 1 Bent/V-shaped < 120° x x

xx

x x

xxO OSx x

= =

2 2 Bent/V-shaped < 109.5°H H

O

3 0 Trigonal Planar 120°

x x

xx x

x

xx

x x

xx

x x

xx

xx

BF

F F

=

3 1Trigonal

Pyramidal<

109.5°

x x

N HH H

4 0 Tetrahedral 109.5° C HH H

H

STOP AND CHECK:

Turn your book over and see if you can figure out the molecular shapes for H2O, CH4 and O2, using the Lewis Diagrams you drew in the last section.

Quick Questions

What determines the shape of different molecules? Why is important to take into account both bonding and non-bonding pairs of electrons?

Try and recreate the beautifully-designed table for the list of molecular shapes, including the number of bonding and non-bonding electron pairs, the name of the shape, and the size of the bonding angles. Chuck in your own examples if you can to sweeten the deal.

?



POLARITYPolarity is one of those great topics that ALWAYS comes up in the exam, and you are always just asked to explain why a molecule is polar or non-polar (or to determine the polarity of a molecule and explain your reasoning). So, for those E8s to come raining down it would be wise to cover:

What polarity means.Electronegativity - what it means and how it contributes to polarity.How non-polar covalent bonds are formed. How polar covalent bonds are formed. What the exam will ask you! We’ll run through a full-proof method of determining the polarity of molecules (which doesn’t involve covering your eyes and circling an option at random).

Arctic (North Pole)

magnet

Antarctica (South Pole)

S

Npolarity

On opposite ends of the planet, Earth has a North and South pole. If have a look at magnets, we call one end the north pole and the other end the south pole. Because of this, both the Earth and magnets are said to be polar, or have polarity.

Chemical bonds and molecules can also be polar (or non-polar)

Here, polarity just means the “separation of charge”. Polarity can apply to both molecules and individual bonds - and it’s important to know the difference between them.

Polar molecules will have a positively-charged region and a negatively-charged region, while non-polar molecules will have no real charge difference across the molecule.

Ethane, a nonpolar molecule

No charge difference

Water, a polar molecule

++

-

+ −

Defining Polarity



In a polar covalent bond, one of the atoms ends up with a partial positive charge while the other a partial negative charge. In a non-polar covalent bond, there is no difference in charge.

- Nonpolar covalent bond- Bonding electrons shared equally between two atoms.- No charges on atoms.

- Polar covalent bond- Bonding electrons shared unequally between two atoms.- Partial charges on atoms.

Cl Cl H Clδ+ δ-

STOP AND CHECK:


The definition of polarity when it comes to chemistry. How polarity can be applied to molecules and covalent bonds.


Electronegativity

Before we jump into the deep end of this topic and start calling molecules polar or non-polar, let’s take a step back and think about one of the most important terms in this topic, “electronegativity”.

We know that atoms are attracted to bonding electrons, but it turns out some are more attracted to them than others.

Heʼs so strong

I am falling for him

The stronger the feelings an atom has for electrons, the stronger the attraction, and the tighter the atom will hold onto electrons in a covalent bond.

If it were up to us we would just call this property “electron lust”, but unfortunately a Swedish chemist, Berzelius, beat us to it and gave it the serious name, “electronegativity”.



Electronegativity describes the tendency of an atom to attract bonding electrons

Pretty much, it just means, the more electronegative you are, the better you are at attracting electrons.

So, how do you know which atoms are more attracted to bonding electrons than others?

If you go down a group in the Periodic Table the electronegativity decreases, but if you go across a period (from left to right), the electronegativity increases.

nitrogen

14.007N

7

helium

He4.0026

2

neon

Ne20.180

10fluorine

F18.998

9oxygen

O15.999

8carbon

C12.011

6boron

B10.811

5

argon

Ar39.948

18chlorine

Cl35.453

17sulfur

S32.065

16phosphorus

P30.974

15silicon

Si28.086

14aluminium

Al26.982

13

krypton

Kr83.798

36bromine

Br79.904

35selenium

Se78.96

34arsenic

As74.922

33germanium

Ge72.64

32gallium

Ga69.723

31zinc

Zn65.38

30copper

Cu63.546

29nickel

Ni58.693

28cobalt

Co58.933

27iron

Fe55.845

26manganese

Mn54.938

25chromium

Cr51.996

24vanadium

V50.942

23titanium

Ti47.867

22scandium

Sc44.956

21calcium

Ca40.078

20potassium

K39.098

19

magnesium

Mg24.305

12sodium

Na22.990

11

beryllium

Be9.0122

4lithium

Li6.941

3

hydrogen

H1.0079

1

1 18

2

3 4 5 6 7 8 9 10 11 12

13 14 15 16 17

xenon

Xe131.29

54iodine

I126.90

53tellurium

Te127.60

52antimony

Sb121.76

51tin

Sn118.71

50indium

In114.82

49cadmium

Cd112.41

48silver

Ag107.87

47palladium

Pd106.42

46rhodium

Rh102.91

45ruthenium

Ru101.07

44technetium

Tc[98]

43molybdenum

Mo95.96

42niobium

Nb92.906

41zirconium

Zr91.224

40yttrium

Y88.906

39

Lanthanite

57-71

strontium

Sr87.62

38rubidium

Rb85.468

37

radon

Rn[222]

86astatine

At[210]

85polonium

Po[209]

84bismuth

Bi208.98

83lead

Pb207.2

82

dysprosium

Dy162.50

66terbium

Tb158.93

65gadolinium

Gd157.25

64europium

Eu151.96

63samarium

Sm150.36

62promethium

Pm[145]

61neodymium

Nd144.24

60praseodymium

Pr140.91

59cerium

Ce140.12

58lanthanum

La138.91

57

barium

Ba137.33

56caesium

Cs132.91

55

roentgenium

Rg[272]

111darmstadtium

Ds[271]

110meitnerium

Mt[268]

109hassium

Hs[277]

108bohrium

Bh[264]

107seaborgium

Sg[266]

106dubnium

Db[262]

105rutherfordium

Rf[261]

104radium

Ra[226]

88francium

Fr[223]

87

lutetium

Lu174.97

71ytterbium

Yb173.05

70thulium

Tm168.93

69erbium

Er167.26

68holmium

Ho164.93

67

thallium

Tl204.38

81mercury

Hg200.59

80

ununoctium

Uuo294

118ununseptium

Uus294

117livermorium

Lv293

116ununpentium

Uup288

115flerovium

Fl289

114ununtrum

Uut284

113copernicium

Cn277

112

gold

Au196.97

79platinum

Pt195.08

78iridium

Ir192.22

77osmium

Os190.23

76rhenium

Re186.21

75tungsten

W183.84

74tantalum

Ta180.95

73hafnium

Hf178.49

72

berkelium

Bk[247]

97lawrencium

Lr[262]

103nobelium

No[259]

102mendelevium

Md[258]

101fermium

Fm[257]

100einsteinium

Es[252]

99californium

Cf[251]

98curium

Cm[247]

96americium

Am[243]

95plutonium

Pu[244]

94neptunium

Np[237]

93uranium

U238.03

92protactinium

Pa231.04

91thorium

Th232.04

90actinium

Ac[227]

89

89-103

Actinide

increasing electronegativity

incr

easin

g el

ectro

nega

tivity

increasing electro

negativity

This means that the least electronegative atoms are cesium and francium, whereas the most electronegative atoms are nitrogen, oxygen and fluorine. (We ignore the Noble Gases since they are unreactive).

For now, just remember that the most electronegative atoms are F, O, N and Cl.

STOP AND CHECK:


What electronegativity is. Which atoms have higher electronegativity values, and which ones have

lower values.




Non-polar Covalent Bonds

Let’s begin this section by grabbing two identical non-metal atoms, say two hydrogen atoms

They both have 1 valence electron but need 2 in total to be stable. Sharing them sounds like a great idea at this point!

HHH H

Even though they want to share, both hydrogen atoms secretly want the electrons to themselves. So, there is a bit of tug of war going on. But, because both atoms are the same they pull on these bonding electrons with the same amount of strength. This means that bonding electrons will happily zip around the nucleus of both hydrogen atoms, spending the same amount of time around each one.

This means there is no real difference in charge between these two atoms, so their covalent bond is non-polar.

A non-polar bond is easy to picture, because it will always be symmetrical!

symetrical electrondistribution

H He

e

STOP AND CHECK:


What are the requirements to make a non-polar covalent bond in terms of the atoms – are they the same or different atoms?

Why non-polar covalent bonds have no separation of charge.




Polar Covalent Bonds

Let’s see what happens when we get two different non-metal atoms, say a hydrogen atom and a chlorine atom

Chlorine has 7 valence electrons but needs 8, while hydrogen has 1 but needs 2. So, they both decide to share 1 electron.

Hydrogen atomshared pair

Chloride atom

HCl

Hydrogen Chloride HCl

HCl

Since chlorine is more electronegative than hydrogen (remember FONCL), chlorine has a larger tendency to attract the bonding electrons. This means that bonding electrons are attracted more to the chlorine than they are to the hydrogen.

e- e-

Although the 2 bonding electrons are shared, chlorine pulls more tightly on them. We end up with the bonding electrons spending more time around chlorine nucleus than with hydrogen.

H Clδ+ δ-

Partial Charges:

Since chlorine isn’t being fair and is being greedy instead (and because electrons are negatively charged), it ends up being slightly more negative than Hydrogen.



When this happens, we say that Chlorine has a partial negative charge, which is represented by the “delta negative” symbol (δ-).

As this happens, it makes sense that Hydrogen ends up with a partial positive charge, which is represented by the “delta positive” symbol (δ+).

This kind of covalent bond is polar

That’s because there is a separation of electrical charge. Since it is polar it creates what is called a, “dipole”.

The letters ‘di-’ mean ‘two’ and ‘pole’ means one end - so a dipole is a way of describing the separation of electrical charge that we just talked about.

uneven electrondistribution

partial negativechargepartial positive

chargeH Cl

e

eδ+ δ+

STOP AND CHECK:


What are the requirements to make a polar covalent bond in terms of the atoms – are they the same or different atoms?

Why polar covalent bonds have no separation of charge.


Polarity of Molecules

With all that sorted, and an understanding of how polarity works on the bond level, let’s have a look at an entire molecule:

Let’s start things off simply by assuming there is only one type of atom around the central atom in the molecule, like in methane (CH4):

H

C HH

H



Here, the polarity of the molecule depends upon the symmetry of the molecule’s shape

Remember the shapes we spoke about in the last section. Well, we promised they would come in handy! It turns out some shapes are symmetrical, whilst some aren’t.

To work out whether a shape is symmetrical or not, picture pulling on each of the atoms surrounding the central atom. If you think there would be movement, chances are the molecule is asymmetrical.

O

HH

Imagine Oxygenʻpullingʼ on both Hydrogens

Overall thereis a ʻpullʼ

Sometimes it can be hard to tell which shapes are symmetrical and which are asymmetrical, so to double check, refer to the table below:

Symmetrical Molecule Shapes Asymmetrical Molecule Shapes

Linear Bent/V-shape

Trigonal planar Trigonal pyramidal

Tetrahedral

Even when the bonds around the central atom are polar, if the molecule has a symmetrical shape the polar bonds cancel out, leaving the molecule non-polar overall. It’s all down to symmetry!

partially -ve

dipoles cancel out

partially +ve

dipole

linearδ - δ+ δ -

C= =O O

C= =O OOn the other hand, when the molecule is asymmetrical, this cancelling of polar bonds doesn’t occur. This leaves the molecule polar instead.



dipoles donʼt cancel out

partially -ve

partially +ve

dipole

bent (V-shaped)

δ -

δ+ δ+

O

HHO

HH

But what about when there is more than one type of atom around the central atom?

In this case the molecule will ALWAYS be polar, regardless of the molecule’s shape and symmetry.

This is because, the different atoms with different electronegativity values produce different strength dipoles. These different dipoles will not be cancelled out even when the molecular shape is symmetrical.

STOP AND CHECK:


What helps us to determine polarity. The importance of symmetry when it comes to polarity. Why molecules with more than one type of atom around the central atom are

always polar no matter what.


Quick Questions What does polarity refer to? What does it mean in context of chemistry? Explain the concept of electronegativity and describe the trends in electronegativity down the columns and across the rows of the Periodic Table.

What is the difference between non-polar and polar covalent bonds? Why is water (H2O) a polar molecule? Why is methane (CH4) non-polar but chloromethane (CH3Cl) polar, even though both molecules are tetrahedral?

?



TYPES OF SOLIDSFor Level 2 Chemistry there are four types of solids that you need to know:

1. Molecular solids2. Ionic solids3. Metallic solids4. Covalent network solids

You will most likely be given an atom or molecular formula and you must be able to describe what type of solid it will be, the type of particles involved and the attractive forces between particles.

We’ll get into it all later, but as a bit of a spoiler, the table below can be used as a quick revision guide for the structure and bonding within each of the different types of solids:

Type of Solid Type of Particles Attractive Forces between Particles

Molecular solid MoleculesWeak intermolecular forces

(Van der Waals forces)

Ionic solid Cations and anions Ionic bonds

Metallic solidMetal cations (and

delocalised electrons)Metallic bonds

Covalent network solid Atoms Covalent bonds

As we learn about each of these types of solids, we’ll be learning how to connect their structures to how they act in different scenarios. These are called their physical properties.

The main properties that are covered for all solid types in this section include:

Melting/boiling pointSolubility Electrical conductivity

For each one you need to know what it is and what factors affect it. Then take what you know about the solid’s structure and bonding to explain what properties it has. We’ll even introduce you to some more properties such as ductility and malleability - which are important when it comes to solids.



Introduction to Solid Properties: Melting/Boiling Point, Solubility and Electrical Conductivity

Before we begin talking about the different types of solids, we should clarify the different properties you’re going to want to know about each one.

As part of this external topic, you must be able to link the bonding and structure of each solid type to the properties it has. The main properties you must be able to discuss are the melting/boiling point, solubility, and electrical conductivity of each solid.

Melting/Boiling Point

All substances can exist as gases, liquids and solids at different temperatures and pressures

For example, water molecules exist as a solid (ice) at temperatures below 0°C. As we turn up the dial and get things heated up, the ice melts and forms liquid water. If we want to take things to the next level, we can boil the water up to 100°C and produce a gaseous form of water (steam).

When the temperature increases, the amount of heat energy increases

This is used to break any bonds holding the solid or liquid together. As a general rule, the stronger the force of attraction the more heat energy is required to break it.

The melting point is the temperature at which solid melts into a liquid, while the boiling point tells us what temperature is needed to boil that liquid into a gas.

solid

intermolecular force

melting

liquid

gas

boiling

Melting and boiling



Solubility:

Solubility tells us how likely something is to dissolve in something else. If a compound is soluble in water it will dissolve (and completely disappear) when added.

The golden rule of solubility is that “like dissolves in like”

This is where the polarity concept we spoke about in the last section becomes super handy.

Let’s start with water. We know that water is a bent shape, so is asymmetrical and hence polar.

O

H H+ +

-

This means that polar molecules can dissolve in water - which is a polar liquid itself, but not in non-polar liquids. You need to know that water is a polar liquid but you will always be told if a solvent is non-polar. For example, in a question it might look like “hexane, a non-polar solvent…”

On the other hand, non-polar molecules dissolve in other non-polar liquids, but not in water.

The thing that gets dissolved in the liquid is called the “solute”, while the liquid it dissolves in, such as water, is called the “solvent”.

water molecule

solid


a solution

bonding with water

O

O

O

O

H

H

H

H

H

H

H

H

dissolving

O

H

H

Solubility



Electrical Conductivity:

If a compound is electrically conductive it can allow a current (the flow of positive or negative charges) to flow through it.

In order to be electrically conductive a substance needs either delocalised, or free moving, charges.

electron

Free to move

fixed positive charge

+ + + + + +

+-

-

-

-

--

+ + + + +

+ + + + + +

+ + + + + +

+ + + + + +

In chemistry, the charges can come as either electrons or in the form of ions (both cations and anions count for this one)!

Keep this in mind as we look at each type of solid as it is important to determine whether it is electrons or ions that are conduct the electrical current.

STOP AND CHECK:


How solids are converted to liquids and how liquids are converted to gases. What solubility means, and the rules that determine if something is soluble in

something else. The requirements for a compound to be electrically conductive.


Molecular Solid Structure

Molecular solids are simply made up of molecules.

Molecules are just a gathering of atoms all sharing their bonding electrons to form covalent bonds.



CC

H

HC C

H

H

H

H H

H H

Molecular solids have a strict, exclusive membership policy, only allowing non-metal atoms to join. The non-metal atoms can be found on the right-side of the Periodic Table. So, if you see something like oxygen (O2), hydrogen chloride (HCl) or ammonia (NH3), you know it must form a molecular solid!

For molecular solids, both intramolecular and intermolecular forces are important

To tell the difference between the two types of forces it might help to think of “intranet” vs “internet”. In this case, think of an individual molecule as a school or company, and the atoms in each molecule as the students or workers.

The intranet is a bit like a private version of the internet that only the students or workers can access, while the internet allows everybody to be connected – it allows schools to connect to other schools and companies to connect with other companies.

Intramolecular forces are private and can only be used within each molecule. These are the covalent bonds that connect the atoms and hold the molecules together.

Intramolecular force (strong)

Intermolecular force (weak)

HClHCl

Intermolecular forces, however, allow each molecule to be connected to other molecules. In molecular solids, these intermolecular forces are called “van der Waals forces”. These intermolecular forces occur between all molecules – whether they are polar or non-polar.

Gases form when each molecule is free to move on its own, but we get liquids and solids when more and more intermolecular forces are formed between molecules.



Least IntermolecularForces

Gas

Solid

Liquid

Most IntermolecularForces

When it comes to strength, intramolecular forces are strong, while intermolecular forces are weak

This means that more energy in the form of heat is required to break covalent bonds, with less heat needed to break van der Waals forces.


intramolecular forcecovalent bond

Molecule H HH

N

H HH

N

H HH

N

STOP AND CHECK:


What makes up molecular solids. How molecular solids are held together.

Try to explain it in your own words, and maybe throw in a diagram.

Molecular Solid Properties

Melting/Boiling Point:

Remember how we said Van der Waals forces are weak? Well, take a beaker of water, crank up the heat, and soon those intermolecular forces are gone. This is because they don’t require much energy to break. As we from solid to liquid the intermolecular forces are somewhat loosened and then when we go to gas phase they are completely gone.



It’s important to be aware that the intramolecular forces never break when we go from solid to liquid to gas.

Compared to other types of solids, the melting/boiling points of molecular compounds tend to be quite low

In fact, most molecular compounds, like methane (CH4), oxygen (O2) and carbon dioxide (CO2), are gases at room temperature.

Not all molecular compounds are the same, and some are more resilient than others. The bigger the molecules are, the better. This is because as the molecular weight of the molecule increases, the strength of the intermolecular forces increases as well. The stronger those bonds are, the more energy (or a higher temperature) is required to break them.

Small molecules Bigger molecules

(low boiling point) (higher boiling point)

intermolecularforce

Solubility:

Remember the golden rule: “like dissolves in like”.

The solubility of molecular solids in water depends on whether the molecules are polar or non-polar

Since water is a polar molecule, only other polar molecules will be able to be dissolved in it.

A non-polar molecule just won’t mix with water. Instead, what will happen is that two separate layers form: a water layer on the bottom and the non-polar liquid on top. Non-polar molecules will only be soluble in other non-polar compounds.

water

Oil



Electrical Conductivity

When it comes to electrical conductivity, molecular solids don’t have what it takes!

These molecules overall aren’t charged. Even polar molecules, which have some charge separation in their molecule, are still neutral overall.

With no charges, there is no electrical conductivity. Therefore, we can say that all molecular compounds are poor electrical conductors (or simply, non-conductive).

STOP AND CHECK:


The properties of molecular solids. Why molecular solids have the properties they do. How to figure out if a molecular solid is soluble in water.


Ionic Solid Structure

Ionic compounds are composed of at least one metal cation and one non-metal anion.

Flashback to page 2: metal atoms tend to lose their valence electrons and pass them along to non-metal atoms who need to gain electrons to fill their valence shells.

These metal cations and non-metal anions are held together by the attraction between the positive and negative charges; opposites attract. This is also called an ionic bond.

Transfer of electron

Atoms

positive ion negative ion

+

+

-+

-

-

-+

So, we’ve got one ionic bond, but we don’t quite have a structure.



Instead of having separate intramolecular and intermolecular forces, numbers of ions can come together to form an ionic solid with a 3-dimensional structure connected by large amounts of ionic bonds.

We call this very important structure a ‘lattice’.

Ionic Solid

electrostatic attraction

anion (- ve ion)

cation (+ ve ion)

+

+

+

-

-

-

+

-

-

+

+

-

STOP AND CHECK:


What makes up ionic solids. What holds ionic solids together.

Try to explain it in your own words, and maybe even throw in a diagram.

Ionic Solid Properties


Unlike the weak intermolecular forces holding molecular solids together, ionic bonds can handle the heat (literally). They are a lot stronger and require a lot more energy to break, which means ionic solids have very high melting and boiling points. For example, the melting point of NaCl is 801°C.

Solubility:

Ionic solids stand no chance against water.

As soon as they take a splash they dissolve away, the ions becoming separated and free to move in the solution

Sound a bit harsh? Here’s what happens:



Remember how water is polar, and looks a little like this?:

O

H H+ +

-

The partially-negatively-charged oxygen atoms in water grab hold of the positively-charged ions, and the partially-positively-charged hydrogen atoms in water grab hold of the negatively-charged ions.

Cl-Na+

OO

O

H H

H

H

HH

O

H

H

H

H H

H

O O

It comes back to this basic idea of “like dissolves in like”. Ionic solids have charge separation with their positive and negative ions - and so will dissolve in polar liquids, like water.

The attraction between the water molecules and the ions is great enough the overcome the ionic bonds, and the individual ions are pulled apart and surrounded by water.


With no delocalised electrons or free-moving ions, solid ionic compounds are poor electrical conductors

Remember that strong lattice structure we talked about?

In the solid state, all ions are tightly held by the ionic bonds. Because these attractive forces are strong, the valence electrons are locked away, held in the valence shells of the ions.

However, if ionic solids are dissolved or melted, they can conduct electricity

In fact they become good electrical conductors. When they are dissolved or molten, the ionic bonds holding the ions together are broken leaving the individual ions, which



are now free to move. When a current is applied, the ionic solution can conduct this current with the movement of positively- and negatively-charged ions.

+

++ +

+ +

+ + +++

++

+

++

-- ++ -

--

- - - -

-- - -

-- -

---

-+ -- ++ -

+ -- ++ -

+ -- ++ -

+ -- ++ -

+ -- ++ -

Solid

Ions fixed in lattice and cannot move

+

- Molten or aqueous solution

ions can now move and conduct electricy

+ ions move tonegative terminal

- ions move topositive terminal

+

Ionic solids are characteristically brittle and non-ductile

If something is brittle it is unable to bend, instead it will snap or shatter when a force is applied. If something is non-ductile it is not able to be stretched into wires.

The reason for this is that the ionic bonds hold the ions in definite positions. So, when a force is applied the negatively-charged ions move over to a point where they no longer line up to a corresponding positively-charged ion, and instead line up with another negatively-charged ion. As like charges repel, the lattice structure shatters.

- + Stress

StressRepulsion

- +

- + - +- + -+

- + -+

- + - +- + -+

- + - +- + -+

- + - +- + -+

- + - +- + -+

STOP AND CHECK:


The properties of ionic solids. Why ionic solids have the properties they do. Why ionic compounds are conductive in the dissolved (aqueous) or melted

state but not in the solid state.




Metallic Solids

Metallic solids, as the name suggests, are made up of metal atoms only

The metals are found in groups 1 to 12 (apart from hydrogen), as well as aluminium, tin, and lead.

We gain metallic solids through the close packing together of metal atoms.

As this packing together happens, the valence electrons are stripped off, leaving the positively-charged metal nuclei

“Sea” ofDelocalised

ElectronsMetal

Cations

+ + + ++ + + + +

+ + + + ++ + + +

The valence electrons no longer belong to one single metal atom, but move throughout the structure like an ocean or a sea. Because they don’t belong to any specific atom or location, we call these ‘delocalised electrons’.

The abandoned positively charged nuclei are still attracted to the constantly moving sea of electrons. This constant attraction results in a cohesive metallic structure.

The metallic solid structure is therefore said to be “held together by the electrostatic attraction between positively-charged metal nuclei and the negatively-charged delocalised electrons”. This is conveniently called a “metallic bond”.

metalnucleui

e

delocalised valence electrons

electrostaticattraction

+

+

e

e+

+

+

+

+

+

+

+

+

+e e

e

e

STOP AND CHECK:


What makes up metallic solids.



What holds metallic solids together.


Metallic Solid Properties


Just like ionic bonds, metallic bonds are quite strong, which gives metallic solids very high melting/boiling points. For example, copper has a melting point of 1085°C.

Solubility:

However, unlike ionic solids, metallic solids are very insoluble in water.

This is because the metallic bonds are unable to be overcome by the water molecules. What this means, is that the water molecules are unable to form strong enough bonds to the positively-charged metal nuclei to pull the metallic atoms apart.


Remember how electrical conductivity is all about free moving charges?

Well, it turns out all metals are good electrical conductors. As the delocalised electrons are free to move within the metallic solids structure, they are able to flow as part of an electrical current.

Metal nuclei

electrons can flow

+ + + + +e e e e e e+ + + + +

e e e e e+ + + + +

e e e e e+ + + + +

e e e e e+ + + + +

e e e e e

e

e

e

e

e

e

e

Metals are also known for their malleability and ductility.

Being malleable means that the solid is able to be bent into different shapes or pressed into thin sheets, while being ductile means that the solid is able to be stretched into a wire.

The metallic bond is said to be non-directional, which means that the attraction between positively-charged metal nuclei and negatively-charged delocalised electrons have no



specific direction, and can therefore occur in any direction. This is unlike a covalent bond, which is directional.

When the metallic solid is put under pressure, the shape can change because the metallic bond is still able to form and remain intact. So, the metallic bond allows metals to be malleable and ductile.

Metal nuclei

ForceLayers ofatoms slide

Force

STOP AND CHECK:


The properties of metallic solids. Why metallic solids have the properties they do.


Covalent Network Solids

As the name suggests, covalent network solids involve covalent bonds. However, they differ from molecular solids as they don’t have intermolecular forces.

Instead, covalent network solids are all composed of atoms each covalently bonded to one another to form a network of billions and billions of atoms. There are three examples that you need to be familiar with: diamond, graphite and silicon dioxide.

DiamondDiamond is made entirely of carbon atoms, where each carbon atom is covalently bonded to four other carbon atoms.

covalent bondcarbon atom

COMPLETE UNIT

middle carbonbonded to 4 carbon atoms

C

C C

C

C

C

C

C

C

CC

C



Because covalent bonds are strong intramolecular forces, the melting and boiling point of diamond is very high! It takes a large amount of energy to break these bonds between the individual carbon atoms. Remember, as a covalent NETWORK solid, the structure goes on forever. There are none of the weak intermolecular forces present and only the strong intramolecular forces of the covalent bonds.

Similarly, the covalent bonds in diamond are much too strong to be overcome by water molecules, making diamond very insoluble in water.

With each carbon’s valence electrons held tightly within these covalent bonds, they are unable to move as part of an electrical current, making diamond not a conductor of electricity.

STOP AND CHECK:


What makes up diamond. What holds diamond together. The properties of diamond. Why diamond has the properties it does.

Try explain it in your own words, and maybe even throw in a diagram.

Graphite

Graphite is also made entirely of carbon atoms; however, each carbon atom is covalently bonded to only three other carbon atoms.

In fact, graphite is composed of layers upon layers of 2-dimensional sheets of carbon

Within each sheet, each carbon atom is covalently bonded to just three other carbon atoms.

Each 2D carbon sheet is then held together by weak intermolecular forces. Since the forces are weak the 2D sheets can be rubbed off from each other, making graphite a very useful material, particularly in pencils. The markings left by graphite pencils are actually the 2D carbon sheets of graphite rubbing off onto the paper.

Carbon atoms have four valence electrons, so in graphite each carbon atom has 1 non-bonding valence electron.

Similar to in metallic solids, these ‘spare’ electrons float between the carbon sheets as a ‘sea of electrons’.



covalent bond

2D carbon sheet

C

CC

carbon atom

van der Waals forces(intermolecular forces)

CC

C

C

C

C

C

C

C

C

C

C

C

C

CC

CC

C

C

C

C

C

C

C

C

C

C

C

e

e e

Like in metallic solids, the presence of these delocalised valence electrons makes graphite electrically conductive

When it comes to melting and boiling graphite, a high temperature is required.

Although the attractive forces between each 2D carbon sheet are weak, the individual carbon atoms within each sheet are still held together by the very strong covalent bonds.

Just like diamond, graphite is also insoluble in water

STOP AND CHECK:


What makes up graphite. What holds graphite together. The properties of graphite. Why graphite has the properties it does.


Silicon Dioxide (Silica)

The last covalent network solid that you need to know is silicon dioxide, or silica. Silicon dioxide (SiO2) is made of silicon atoms each covalently bonded to four oxygen atoms, and each oxygen atom bonded to two silicon atoms, resulting in a similar geometric structure to diamond - just with different atoms involved.

Luckily, we already know the properties here! Just like diamond, silicon dioxide has a very high melting and boiling point, is very insoluble in water, and is not a conductor of electricity.

Graphite



OSi

STOP AND CHECK:


What makes up silica. What holds silica together. The properties of silica. Why silica has the properties it does.

Try explain it in your own words, and maybe even throw in a diagram.

Quick Questions Match the following with the type of solid that they will form: water (H2O),

magnesium (Mg), aluminium oxide (Al2O3), and diamond. For each one, describe the structure and bonding that is present.

Compare the melting/boiling point, solubility in water, and electrical conductivity of each of the four compounds above.

Why is diamond a poor electrical conductor but graphite is a good electrical conductor, even though both are covalent network solids made entirely of carbon atoms.

Compare the malleability and ductility of metallic solids and ionic solids.

?



THERMOCHEMISTRYThere are so many chemical reactions that can occur, with almost endless different products formed from heaps of different reactants. But they all have two things in common:

1. They all involve bonds breaking and/or bonds forming.2. They all involve a change in energy.

From just these two points this topic blows up and covers:

Enthalpy and enthalpy change, featuring a few formulae that you have to memorise Endothermic and exothermic reactionsActivation energy

Enthalpy and Enthalpy Change

The first thing to remember is that:

All molecules possess a specific amount of energy

This specific amount of energy is referred to by a special term called ‘enthalpy’, which is stored within the chemical bonds inside the molecule.

Energy is required to break these bonds and energy is released when these bonds are formed.

Enthalpy

Breaking bonds(takes in energy)

Making new bonds(gives out energy)

C

C

C

H

+

H

H H

H H H

HH

HO O

O O

O O O O

O O

O

H H

O

Therefore, in a chemical reaction, the energy of the reactants will be different to the energy of the products, as bonds are broken and formed to make different molecules - each with different amounts of energy.



This change in energy is referred to as the enthalpy change of reaction

One way to calculate the enthalpy change of a reaction is to take the energy required to break all bonds and subtract the energy released when all bonds are formed. The enthalpy change of a reaction has the symbol, ΔrH. Let’s break this symbol down:

The triangle, “delta”, means “change in”. Remember, we’re looking at the enthalpy of a reaction, it is the enthalpy change of a reaction. The small “r” just stands for “reaction”. Finally, “H” is the symbol for enthalpy.

Remember, energy is conserved and cannot be created or destroyed.

If energy is required by the chemical reaction it needs to come from somewhere else, and if energy is released by the reaction it needs to go somewhere.

The actual chemical reaction, containing all the reactants and products, is called the “system”. The “surroundings” is everything that’s not the system; the rest of the universe!

water molecules: water, iceSystem

everything else: cup, air, etc.Surroundings

Energy can come from the surroundings or energy can be released into the surroundings. If you get these two terms nailed down, try use them when discussing enthalpy changes for maximum kudos.

STOP AND CHECK:


The definition of enthalpy. What is meant by the “system” and the “surroundings”.




Endothermic Reactions

There are two types of enthalpy change: endothermic reactions and exothermic reactions

The best thing about being a scientist is getting to use lots of complicated words; it’s like learning a new language. The better news, is that these words often give use clues to what we are talking about. For example, “endo-“ means “in” and “thermic” refers to “heat”, which makes “endothermic” just “in-heat”.

This is great for us because it tells us that:

In an endothermic reaction heat energy goes into the system from the surroundings

If we take away heat from the surroundings, things get a tad chilly and the temperature drops. Often heat and energy is used interchangeably in this standard and that’s because heat is just a form of energy that particles and other objects have which can flow from one to another due to temperature differences.

Looking at the system, because they have drawn in energy, the products end up with more energy than the reactants, meaning endothermic reactions have a positive enthalpy change.

Picture yourself holding a cube of ice

Your hand starts to get really cold as the ice begins to melt. This is because the ice cube acts like the Dementors from Harry Potter, but rather than sucking out your soul it sucks out your heat.

The heat leaves your hand (the surroundings) and enters the ice cube (the system), making this an example of an endothermic reaction. This heat energy is then used to heat up the ice and then used to break the intermolecular forces between water molecules to convert it from a solid to a liquid.



Energy Profile Diagram for an Endothermic Reaction

An energy profile diagram simply tells us how much energy is present in a reaction as time passes.

The one thing to look for in these types of diagrams is whether the products have more or less energy than the reactants. This is important, as it tells us whether heat energy entered or exited the system.

In this case, because the reactants have more energy than the products, we know that energy (in the form of heat) must have entered the system:

reaction time

ENERGY

activationenergy products

reactants

H

Don’t worry about what activation energy in the diagram is just yet - we’ll cover that soon!

STOP AND CHECK:


The energy diagram for an endothermic reaction – make sure to include the activation energy and all labels!

The energy transfer that takes place in an endothermic reaction – where is energy absorbed from and where does it enter?




Exothermic Reactions

If endothermic means “in-heat”, you might guess that exothermic basically means “out-heat”.

So, in an exothermic reaction we know that heat energy must go out of the system and go into the surroundings

This means that the products end up having less energy that the reactants, giving us a negative enthalpy change.

As the heat in the surroundings increases the temperature rises.

Exothermic reactions are the most exciting in chemistry as they include all the explosions and fire that any normal person loves.

For example, combustion reactions (which involve burning stuff) are a classic example of an exothermic reaction

We start off with some crazy, or unstable, reactants (fuel and oxygen), and end up with boring, stable, products (carbon dioxide and water).

Nothing gets a molecule hyped up like a bit of heat energy. So, to calm it down, heat energy needs to be removed from the system. It must go somewhere, so it’s released into the surroundings.

The heat energy released causes a rise in the surrounding temperature, and in a combustion reaction the intense heat released produces flames!



Energy Profile Diagram for an Exothermic Reaction

The energy profile diagram for an exothermic reaction looks like this:

activationenergyreactants

reaction time

ENERGY

productsreactants

H

Don’t worry about what activation energy in the diagram is just yet - we’ll cover that soon!

STOP AND CHECK:


The energy diagram for an exothermic reaction – make sure to include the activation energy and all labels!

The energy transfer that takes place in an exothermic reaction – where is energy absorbed from and where does it enter?


Activation Energy

You may have noticed that in each of these energy profile diagrams there is a small bump in energy levels just after the reaction starts.

These small bumps are the activation energy of a reaction.

This activation energy is present in both endothermic and exothermic reactions.

Activation energy is formally defined as, “the energy barrier that is needed to be overcome before a reaction (between the particles) can take place”. Really, it is the “spark”, or “push”, that gets all reactions going.

The amount of energy required is specific for every reaction, with some reactions only having a small amount of activation energy, while others having a very high activation energy.



STOP AND CHECK:


The definition of activation energy, and why it is important in reactions. How the size of activation energy differs in an endothermic and exothermic

reaction.


Enthalpy Calculations Part 1The enthalpy change of reaction, which is measured in kilojoules per mole, or kJ mol-1, is useful as it allows us to calculate the amount of heat released by a particular amount of substance. This involves multiplying the ΔrH by the number of moles used in the reaction.

Before we can do that we need to know the ΔrH and the number of moles being reacted.

The formula for calculating the number of moles from the mass of a substance is:

n = m ÷ M

In this formula, ‘n’ is the number of moles of a substance, ‘m’ is the mass of a substance (measured in grams, g), and M is the molar mass of a substance (measured in grams per mole, g mol-1).

Let’s look into these a bit more:

The ‘mole’ is a unit for the amount of substance

It’s basically a number, a huuuuge number in fact.

One mole of water molecules is equal to 6.022 x 1023 molecules, one mole of hydrogen atoms is equal to 6.022 x 1023 atoms, and one mole of oxide (O2-) ions is equal to 6.022 x 1023 ions.

It seems like a bit of a random number, but it’s based around the number of atoms in 12g of carbon-12.

We use moles in our calculations as they are a handy way of expressing very large numbers in numbers we find a lot more friendly. For now, just think of a mole as a specific quantity, like a ‘dozen’.



The molar mass measures the size of each mole we are dealing with.

The molar mass can be calculated by adding the atomic weights of all the atoms in the molecule

Just like humans have a weight associated to them, so too do atoms.

The atomic weight of an atom can be found on the Periodic Table. If you have a look at one of the elements in the Periodic Table you will see 3 main things:

1. The symbol for the element. 2. Two numbers:

a. The smallest one, which is always a whole number, is the atomic number which tells us the number of protons in each atom of the same element.

b. The larger number is the atomic weight, which is not usually a whole number. It is also the molar mass, measured in grams per mole.

So, once you have the atomic weight of each atom in the molecule, simply add them all up to get the molar mass of the molecule overall.

For example, to find the molar mass of CO2 we find the atomic weight of carbon (C) and each oxygen (O) , and then add them all up.

• C = 12 g/mol• O = 16 g/mol• The molar mass of CO2 = 12 + (2 x 16) = 44 g/mol

Now that we’ve worked out what we mean by ‘moles’ and ‘molar mass’, let’s jump straight into working out the enthalpy of a reaction with an example:

We’ve got a reaction: H2(g) + ½O2(g) → H2O(g) with an enthalpy change (ΔrH) of -242 kJ/mol.

Imagine we want to calculate how much energy is released when 20 grams of hydrogen gas is combusted.

First of all, we need to calculate the number of moles we are dealing with. If we look at the information, we’ve currently got a value for the mass. We also know that the number of moles can be calculated using n = m ÷ M.

This just means that, in order to calculate the number of moles, we need the molar mass.



If we look at the Periodic Table and find hydrogen:

Element namehydrogen

H1.0079

1 Atomic number

Element symbol

Atomic weight

The molar mass is equal to 1 g/mol (we can round 1.008 down to just 1 in this example). This means the molar mass of H2 is 2 g/mol. So, plugging those values into our equation for the number of moles gives us n = 20g ÷ 2g/mol = 10 moles.

Now that we have our number of moles of hydrogen we can go about calculating the energy released. Looking at our information, we know that the enthalpy change of reaction is equal to -242 kJ per mole. Since there is 242 kJ released per 1 mole of hydrogen released, with 10 moles there will 10 x 242 = 2420 kJ.

The negative sign in front of 242 kJ/mol in the enthalpy change of reaction tells us that this energy is released to the surroundings. Therefore, when talking about the amount of energy released we can drop the negative sign.A positive sign in front of the enthalpy change of reaction would tell us that this energy is absorbed.

STOP AND CHECK:


How to calculate the number of moles (n) from the mass (m) and molar mass (M). How to calculate mass (m) from the number of moles (n) and molar mass (M). How to calculate the enthalpy (ΔH) from the number of moles (n) and the

enthalpy change of reaction (ΔrH). How to calculate the enthalpy change of reaction (ΔrH) from the number of

moles (n) and the enthalpy (ΔH).




Enthalpy Calculations Part 2

An alternative way of calculating enthalpy is by looking at bond enthalpies.

The bond enthalpy is the amount of energy required to break that bond

Using the bond enthalpies, the enthalpy change of reaction can be calculated using the equation: bonds broken - bonds formed.

But what is the “bonds formed” and what is the “bonds broken”? In any reaction we are going from reactants to products. Therefore, to calculate the enthalpy change in reaction we take the enthalpy of the products (the “bonds formed”) and minus the enthalpy of the reactants (the “bonds broken”).

We can calculate the enthalpy of products and reactants using the bond enthalpies of the molecules involved.

Let’s have a look at an example

Say we have the reaction: H2(g) + ½O2(g) → H2O(g). This is the combustion of hydrogen gas. Now, imagine we want to calculate the enthalpy change of reaction ΔrH using bond enthalpies. In questions like these the bond enthalpies will be given to you as a table like this:

Bond Average Bond Enthalpy (kJ/mol)

H-H 436O=O 498O-H 464

To calculate the enthalpy of products and reactants we look at the molecules on each side of the equation and work out how much energy would be required to break the bonds holding the atoms together.

For the products:

H2O is made up of 2 O-H bonds. The bond enthalpy of O-H is 464, and so with 2 O-H bonds it will take 2 x 464 = 928 kJ/mol. Enthalpy of products is 928 kJ/mol.



For the reactants:

We have both hydrogen (H2) and oxygen (O2) to deal with. Let’s look at hydrogen first:

• We have 1 mole of H2, and in each molecule there is 1 H-H bond. • Each H-H bond has an enthalpy of 436 kJ/mol.

How about oxygen: • We have ½ mole of oxygen, and in each molecule there is 1 O=O bond. • Since each O=O bond has an enthalpy of 498 kJ/mol we need ½ x 498 =

249 kJ/mol of energy to break the oxygen molecules apart. Putting it together, the enthalpy of the reactants is 436 + 249 = 685 kJ/mol

Now that we have our “bonds formed” and “bonds broken” we can calculate our enthalpy change of reaction for the combustion of hydrogen:

ΔrH = enthalpy of reactants - enthalpy of products = 685 - 928 = -243 kJ/mol

It is possible to go backwards to calculate the bond enthalpy from the enthalpy change of reaction when other bond enthalpies are known

STOP AND CHECK:


How to calculate the enthalpy change of reaction (ΔrH) using bond enthalpies. What a bond enthalpy is.


Quick Questions What is the “enthalpy change of a reaction” (ΔrH)? Compare exothermic and endothermic reactions, including an energy diagram for each one, showing the transfer of energy between system and surroundings. What is activation energy and why is it important?

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KEY TERMSActivation Energy:

The energy that is needed to be overcome before a chemical reaction can start, or take place.

Bond Enthalpy: The energy required to break a particular chemical bond.

Brittle: A property of a substance whereby the structure snaps, or shatters, when a force is applied.

Covalent Bond: The bond between two atoms that forms when a pair of electrons that are shared between two atoms.

Delocalised Electrons: Electrons that are not assigned, or localised, to one specific atom, but rather are free to move around the structure.

Dipole: A term used to describe a separation of positive and negative charges. For example, a polar bond creates a dipole.

Ductility: The ability of a substance to be stretched (into wires) when a force is applied.

Electronegativity: The tendency of an atom to attract electrons.

Endothermic Reactions: Chemical reactions which involve a transfer, or absorption, of energy from the surroundings into the system.

Enthalpy of Reaction, ΔrH: The enthalpy change associated with a specific chemical reaction and the mole ratios involved. It is the difference between the total enthalpy of the products and the total enthalpy of the reactants when the amount of reactants and products is equal to the mole ratios.

Enthalpy: The particular energy that a molecule possesses, which is determined by the strength of the bonds.



Exothermic Reactions: Chemical reactions which involve a transfer, or release, of energy from the system into the surroundings.

Intermolecular Forces: The attractive forces that act between molecules, holding them together.

Intramolecular Forces: The attractive forces that act within individual molecules, holding the atoms that make up that molecule together.

Ionic Bond: The electrostatic attraction that occurs between positively-charged ions (cations) and negatively-charged ions (anions).

Malleability: The ability of a substance to be bent into different shapes when a force is applied. Often substances that are malleable can be hammered or rolled into thin sheets.

Metallic Bond: The electrostatic attraction between the positively-charged metal nuclei and the negatively-charged delocalised electrons.

Soluble: If a substance can be dissolved in a liquid, called the solvent, then it is said to be soluble in that liquid.

Surroundings:In a chemical reaction, the “surroundings” simply refer to everything else that’s not the system. We can think of this simply as the air around where the reaction is taking place.

System: In a chemical reaction, the “system” refers to all the energy and matter (the particles, or the molecules) involved in a chemical reaction.

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