CHEMICAL BONDS

Introduction

You know that all substances are built up with particles. The most simple particles are atoms or ions. (thats where we stay for now)
Those particles - normally - do nog stay alone, but connect to other particles.
That can happen in different ways, but they always approach and connect at their exterior, their outside; so you can imagine that this must involve the outer electron shell, orbital and the valency electrons in that shell.
Only the atoms of noble gases can remain 'bachelor'. They normally do not connect to other particles.

Every way to connect particles we call that a chemical bond. That is the main topic of this module.
The character of such chemical bonds has big consequences for the properties of the substance.

content:

1. the ionic bond

1.1 the ionic lattice

2. the metallic bond

2.1 the metallic lattice

3. the covalent bond

3.1 molecules

3.2 the polar covalent bond

3.3 ΔE

3.4 double or triple bonds

3.5 aromatic bonds

4. combinations of chemical bonds

5. the chemical bond and energy

6. Intermolecular forces

6.1 vanderWaals forces

6.2 dipole forces

6.3 hydrogen bridges

6.4 molecular lattice

1. The ionic bond

We know simple and complex ions. Both can participate in an ionic bond.
As soon as ions attract each other - because they have opposit charges - you can get an ionic bond.

The bigger purple spheres represent, for example, the negative ions, anbd the grey smaller spheres are the positive metal ions.

You know: there are two kinds of ions: negative and positive ones.
They always are charged, otherwise they wouldn't be ions.

In general, the metal ions will make positive ions (metal atoms like to donate electrons), and the non metal atoms form negative ions (because non metal ions prefer to gain electrons).
In other words: if a metal and a non metal react, electrons are transferred from the metal to the non metal.
That's how positive and negative ions are made; they will immediately attract each other.
Don't forget: the properties of a neutral atom are completely different from the properties of the according ion.
For example: Na, Cl and F in neutral situation are poisonous, but as ions they are found in kitchen salt and in fluorine tablets.

The IONIC bond = attraction between positive and negative particles.

Opposite charges attract each other; equal charges to repell

The force of repulsion or attraction depends on:

how big the charges are
how close the charges are

For this we apply the formula of Physics (there is the Law of Coulomb again):

k = a constant
Q₁ and Q₂ are the charges of the ions
n = the distance (measured from center point to center point; look in a table with ionic rays.

Question 1
Explain that a Chlorine atom will attract electrons more fanatically than an Iondine ion will do.
Don't forget here the concept of electronegativity.

You must always mind: if a neutral substance contains positive ions, there must be negative ions (same total charge) in that substance.
A metal in a compound has before reacted and donated its valency electrons.
That's how you got a positieve charged metal particle.
The valency electrons were transferred to the non metal part, that, of course, became negative.

Another way to say the same:
If you have a connection, a bonding, wherein metals and non metals are present, the metal part has always donated electrons, and is thus positive; The rest of that compound must be negative.
If you see that a chemical compound, a substance, contains a metal, than there is an ionic bond. [an exception may be a compound with an ammonia ion]

If we write ions separately, the charge always must be indicated right above, also if complex ions are involved.
If a substance has a charge of 0, so if there is equal amount of + and -, than don't write any charge in the total formula.

Examples:
Na⁺ + Cl^-

NaCl
I₂ + I^-

I₃^-

Question 2
Explain why the name 'fluorine tablet' in fact is not a good one; what could be a better name?

If non metals have made negative ions, their name will change with a suffix: you add: -ide.

Examples:
Cl^- = chloride; O^2- = oxyde S^2- = sulfide N^3- = nitride

Question 3
Write the formulas of Bromide, Fosfide and sulfide.

Positive ions do not get such a name change.
In the case of a metal with various valencies (ex: Iron ions can be 2+ or 3+) than you must indicate with a Roman number the right valency in the name of the ion:
iron(II)chloride = FeCl₂ (ions Fe²⁺ and Cl^-) The Iron has here a valency of 2
lead(IV)oxyde = PbO₂ (ions: Pb⁴⁺ and O^2-) The Lead has here a valency of 4

A metal is a free element, red Copper for example, or Tin, did not react yet; the atoms still have their valency electrons and still are neutral.

The bigger purple spheres represent the negative ions, and the grey smaller spheres are the positive metal ions.

1.1 The ionic lattice

Substances built up of ions form IONIC LATTICES.
The different ions will pack neatly, in a very regular way.
positive ions are surrounded by negative ions, and negative ions are surrounded by positive ones. Alle because of the attraction forces of course).
Those lattices are strong, because of all those attraction forces.
It is good to know already that most substances with an ionic lattice are called: salts.

Question 4
Try to imagine the ionic lattice of NACl, where the positive and the negative ions are neatly arranged and packed close toghether (see image above).
The values of the ionic charges can be found in tables, just like the values of the ionic rays (so how big the ions are).
Calculate the attraction force between the Na⁺ and Cl^- ions in the cristal of kitchen salt.

Question 5
Indicate for the following compounds if they will have an ionic bond; if yes, than note the ionic formula:

NaCl C₆H₁₂O₆ FeCl₃ Na₂SO₄

All substances built up of ions, do not form molecules.
Molecules are treated further in this module, but it is good to know already that molecules are built up of connected atoms, not as ions.
Molecules are neutral particles.

Also substances with ionic bonds are neutral, but only because there are equal amounts of positive and negative charges.
Positive and negative ions always join toghether in such a way that in total a neutral substance is made: with the same positive and negative charge.

The positive and negative ions join in a certain fixed proportion in unbelievable large amounts, and are very nice arranged toghether in the IONIC LATTICE.

The proportion of pos and neg ions is indicated with numbers right under, as if it were a molecular formula. But it is not a molecular formula. this is a Empiric Formula, of a Proportion Formula.
Att.: It looks as if there is no difference between proportion formulas and molecular formulas.

2. The metalic bond

Metals have just metal atoms. And those atoms normally want to donate electrons to form positive ions.
But: if any substance (like a metal) only has atoms wishing to donate electrons, to where can those valency electrons go? How about that?
Well, it appears that those donated (valency) electrons arrive in a kind of autonomous situation; they join, they move toghether, they stick toghether, and achieve in this way something they appreciate: space.
Besides: this is not only the case in metalic bonds, but in all chemical bonds. Electrons tend to have more space to move in
Better to say: they achieve more stability.

Those valency electrons of the metal atoms, in fact, are belonging to alle metal atoms within the substance; they all move around the atoms.

At the same time you can see it this way: They have in fact donated their valency electrons and became positive ions.

2.1 The metalic lattice

The metal atoms, and also their ions (after losing their val electrons) have their fixed position in a lattice; they are very neatly packed and ordered, just as in an ionic lattice. But now we speak of a METALLIC LATTICE.
All those positive ions are kept toghether by the "free moving electrons".
Those (negative) electrons can be considered as a kind of glue that holds the atom ions, the metal ions, toghether in the lattice.
Those valency electrons are called 'free' because they are no longer connected to only one atom.
The attraction between the negarive free moving electrons and the positive ions in the lattice is what we call: METALLIC BOND.

Metallic bond = attraction between (negative free) electrons and (positive) metal ions in a lattice

These attraction forces can be considerably strong. But can differ per metal.
To break a lattice of Iron (i.e. to melt Iron) you need very high temperatures, but the metal Lead melt much easier. (see module 5)

Question 6 is a practical task
[a practical task must be done in groups.]

You must try to resolve a number (8) of tasks according to a pattern of (ten) questions (question list).
First you have an example (reaction of Iron with Sulphur). With that example we discuss all the questions. So you know what to do at all 8 tasks.
This time you don't need a lab. Its all theory, paper work.

1. Potassium with Chlorine	2. Magnesium with Oxygen	3. Sodium with Sulphur	4. Zincum with Argon
5. Aluminium with Phosphor	6. Silver with Arsenium	7. Cromium with Oxygen	8. Magnesium with Helium

Every time the question is if the two substances do react and, if yes, what you may know about the chemical bonds between the two.
You must study every reaction according to the question list below.

the question list:

Have you seen or done this reaction before? If yes, how was that reaction in practice?

Which formulas do the two elements have? Give also the states (s) (l) (g)

What must happen to both reactants to get loose atoms? Indicate that with formulas

Write the electronic formulas of the two atoms. If needed, use Periodic Table

Are there atoms that want to gain or donate electrons completely? If yes, write the reaction of that process

Is a reaction possible? If not, than stop this one and go to the next reaction

Write the formulas of the product:

ionic formulas of loose ions in the product
proportional formulas and electronic formulas

What type of bonding has been made here?

Write the name of the product. Do you know this substancen in practice?

Make the right reaction equation

Example: Iron with Sulphur

Mr X answering the question list:

I have never seen the reaction of Iron with Sulphur

Fe(s) and S₈(s)

To get loose atoms of Iron, the atoms must get out of the metalic lattice of Iron.
To get loose molecules of Sulphur, the molecules must get out of the molecular lattice of Sulphur; after that, the 8 atoms in one molecule S₈ must separate, get loose from each other.

Write the electronic formulas of the two atoms. If needed, use Periodic Table

The Iron has the tendency to donate electrons completely: Fe

Fe + 2^e
S₈(s)

S₈(g) S₈

Yes, a reaction between this metal and non metal is possible

Fe²⁺ and S^2-
FeS

an ionic bond was formed

Iron(II)sulphide. Never seen before!

8Fe(s) + S₈(s)

8FeS(s)

3. The covalent bond

When two non metals react, two atoms come toghether and both atoms prefer to gain electrons (and to become negative).
You know that non metals already have a good number of valency electrons. For them it is easier to gain and come to the number of 8 electrons.
But... you know, if there is not any metal atom nearby, where do the electrons come from?

Yet these non metal atoms both can get noble gas configuration, if they share each others electrons.
To realise that, these atomd must come very close. Even they overlap in a way; very close thus.
This way the non metals can very well react with one another.
Also in this case the valency electrons get an increased space for movement, they become more stable.

But mind: no valency electrons are transferred now from one atom to the other. No positive and negative ions can be formed.
And no ionic bonds are made.

Compare the following images of two substances: CO₂ and SO₂

Covalent bond = sharing of each others electrons

Both atoms want to gain electrons, but there is here a 'win-win-situation' if they share, if they cooperate.

Suppose for a moment that two non metals react and that these two non metals are P and Q.
One atom of P starts to create a shared electron pair with one atom of Q. So that's the start of the reaction.
If one or both of the atoms have not yet reached noble gas configuration (normally 8 electrons in the outer shells, but in the case of Hydrogen 2) a third atom is involved.
This goes on and on until all involved atoms have got noble gas configuration.
There is a rule here: the same kind of atoms prefer not to connect (except C-atoms).
So if Sulfur and Oxygen atoms come toghether, atoms of S and atoms of O will like to connect, better than two S atoms or two O atoms.

Question 7
Try to write, to draw, what happens if loose atoms of Phosphor and of Oxygen com toghether to react.

Question 8
Chose your answer and explain that choice:

If atoms of Hydrogen and of Oxygen meet, than will be formed:

an ionic bond

a polar covalent bond

a non polar covalent bond

no bonding at all

3.1 The Molecule

If two or more atoms are connected in a covalent way, and neutral particles are mad, than those particles must be molecules.
Mostly we deal then with different atoms, but sometimes with the same atoms.

Question 9
Is kitchen salt built up of molecules? Explain your answer.

Something about the threedimensional shape of molecules.
Have a close look at the models below (of molecules):

left could be tetra chloro methane and right di cloro methane.

the cyclic form of glucose

Molecules are built up of two or more connected atoms.
The occupy space and have a threedimensional shape.
You can have extremely simple molecules, like H-H (H₂).
This Hydrogen molecule is very small, linear and symmetric.
But most molecules occupy much more space, are bigger, like glucose.

Question 10
Controll in every structure:

if the total number of valency electrons is right, and
if the atoms succeeded in realising a stabel distribution of electrons.

Goto answer 03-10

If a particle has an unpaired electron, this particle will be very reactive and is called a radical.

The number of atoms of one kind that is present in a molecule, is indicated with a number right under. That's how you get a MOLECULAR FORMULA.

In the images above you see the molecules CO₂ and SO₂. Every molecule is here built up of three connected atoms.

Question 11
Explain if the following formules are molecular formulas or proportional formulas: H₂O NaCl FeS P₄

Very often two non metals can form a molecule in only one way, but sometimes there are different combinations.
If two non metals can form different molecules, than we should adjust the names.
In that case we apply extra rules for nomenclature, or: we use prefixes (see module 4):

mono	1	di	2	tri	3	tetra	4	penta	5
hexa	6	hepta	7	okta	8	nona	9	deka	10

The name of a substance can, according to these rules, become rather complicated, which not always is needed. Some names are abbreviated.
Look at the table below with a couple of examples. Certainly you have heard of carbon monoxyde or carbon dioxyde.

CO	(mono)carbon mon(o)oxyde
CO₂	(mono)carbon dioxyde
P₂O₅	(di)phosphor pent(a)oxyde
P₂O₃	(di)phosphor trioxyde

Question 12
Write the names of the following substances:

N₂O

N₂O₃

CuCl

PbS

SO₂

Gaining or loosing electrons completely is limited to the number of three; 4+ or 4- ions are very rare.
Such a limitation does not exist for sharing of electrons in covalent bondings.
An atom may have four or six shared electrons with another atom.

In the structural formula of a molecule you can indicate all valency electrons as dots.
Then you have god an ELECTRONIC FORMULA. An electronic formula does not need to limit itselve to one atom.

If you indicate all valency electrons with dots in the formula, you can exactly see where the electrons are shared.
In stead of dots, we mostly use dashes (two dots equals one dash). Normally electrons perform in pairs. One dash between two atoms indicates such a shared electron pair.

Question 13
Give examples of every kind of formula:

molecular formula
structural formula
proportional formula
electronic formula
ionic formula

If atoms share each others electons, they form molecules. In the case of an atom group with a charge, we talk about a complex ion.

If you would investigate alle substances in the earth crust, you might draw the conclusion that molecules dominate; there are more molecules than ionic of metalic substances.

3.2 Polar covalent bond, dipole molecules

A bond between two of such atoms is called "polar"

	I	II	III	IV	V	VI	VII	VIII
1	H							He
2	Li	Be	B	C	N	O	F	Ne
3	Na	Mg	Al	Si	P	S	Cl	Ar
4	K	Ca	Ga	Ge	As	Se	Br	Kr
5	Rb	Rb	In	Sn	Sb	Te	I	Xe
6	Cs	Bi	Tl	Pb	Bi	Po	At	Rn
7	Fr	Ra

metals	Cs
metalloids	Po
non-metals	Se

The scheme above is the simple periodic table (limited to the main groups I - VII)

the red zone: the non metals have the tendency to gain electrons and so become negatively charged (big electronegativity) the blue zone: the metals have a tendency to donate electrons and so become positive (small electronegativity)

Question 14
Explain the details of that scheme.

Question 15
What is the most important reason that bonds between atoms of noble gases or with atoms of ions do hardly exist?

3.3 ΔE

Metalic bonds	many metals have a low value of E
Ionic bonds	ΔE: > ±1,6
Covalent bonds - covalent, non polar - covalent, polar	0 < ΔE < ±1,6 0 < ΔE: < ±0,4 ΔE: > ±0,4

Question 16
Try to understand the above scheme, and explain the difference in electronegativity of the two connected atoms

Att.
In natural sciences we use lots of symbols from the Greek alphabeth, like:
The Greek capital Delta Δ: represents the exact difference between two values.
the Greek small lettre delta δ: represents that there is only a slight difference.

In a covalent bond can exist a difference between de E values (electronegativity) of the two participating atoms.
Then there isa ΔE, that means: a calculatable difference between the electronegativities of those two atoms.
There will be some dislocation of the electrons between the atoms, i.e. of the shared electrons.
The side to where those electons shift will become a bit more negative, because electrons are negative.
The other side of the bond, with the atom that has the bigger attraction (the biggest ΔE), will have a bit of a negative charge that we indicate as: (δ^-).
At the other side of the bonding is an atom with the smaller electronegativity. That atom must let go (a bit) the electrons, and that atom will become a bit postive = (δ⁺).

Covalent bonds can have a complete NON POLAR character if there is no δ^- and no δ⁺.

If inside the covalent bond there is a certain difference, if the shared electrons stay not exactly in the middle, if there is a certain disequilibrium of the charges δ⁺ and δ^-, then we call that bond: POLAR.

Let's look at it in another way:
The bigger the difference of electronegativity E of the elements of a bonding, the more ionic will be the character of that bonding between those elements.

Thus a substance like Al₂O₃ has a 70% ionic character and a 30% covalent character.

Question 17
Given the molecules CCl₄ NH₃ H₂O HF.
What can you say about the bondings in these molecules, looking at the differences in E?

Dipoles / dipole molecules

The following image you might have seen before:

Molecules can contain one or more polar bonds, and they might be, or not, a dipole.
That depends on the symmetria of the molecule (consider the examples).

Examples:

CS₂ (ΔE = ±0)	CO₂ (ΔE = ±1.0)	H₂O (ΔE = ±1.3)
Covalent bondings	Covalent bondings	Covalent bondings
Non polar molecules	Non polar molecules	Polar molecules
There is no dipole	There is no dipole	There is a dipole
There are no δ⁺ and δ^-	There are δ⁺ and δ^- of which the central points do overlap	The central points of δ⁺ and δ^- do not overlap each other (remain at a distance)
S=C=S	O = C = O δ^- δ⁺ δ^-	δ⁺ δ⁺ H H \ / O δ^-

Question 18

Given the molecules of CCl₄ and CH₂Cl₂, (see image) and of NH₃, H₂O, HF
What conclusions can be drawn comparing these molecules in:

Polarity
Dipoles
The ionic character

In general, polar substances will attract each other, and dissolve in polar solvents (e.g. sugar in water).
The same way we can say that non polar substances do also attract each other and dissolve in each other (like fat in oil).
Polar and non polar substances do not mix.
More about this in module 5.

Question 19
A water molecule has no linear, but a triangle structure. The three atoms have not a position in one line. This little difference has an enormous impact in our world. Imagine that the water molecule would be linear. How about the consequences for our world?

Goto answer 03-19

3.4 Double and triple bonds

3-hexene (with a double bond)

Question 20

consult tabel IV and VI and find the next data:
1. table IV: the inter-atomic distances of covalent bonding types
2. table VI: the bonding energies of covalent bondings
Copy from this tables the information about the bonds: C - C, C = C, C ≡ C;
Put those data in a diagramme and make a graph.
The x-axes of this graph shows the three bondings and the y-axes shows the distances as well as the energies.
then you get two different graphs in one scheme.
Look well and draw your conclusions.
Make an estimation of distance and energy of a C ^... C bond in the benzene ring.

Question 21
Is the following statement true?
"In a triple bond, e.g. in acetylene (= ethyne) 2 C-atoms share a total of 6 electrons."
Explain your answer.

3.5 Benzene

The molecular formula of benzene is: C₆H₆

Every Carbon atom in the ring has three normal covalent bonds. That's how every C-atom remains with a 'non paired electron' that still can connect with whatever other atom.
So in total six electrons are in stock, and these six now will connect in a very special way with each other, without other atoms. A special π-bonding (a "pi-bonding) is made. We come tot that later.
Here we do not treat the very complicated bonding types, so you may simpify here by saying:
Those six electrons obtain the opportunity to dislocate all over the ring of the 6 atoms.
They are allowed to move freely over this ring, obtaing this way bigger freedom, more movability.
That means at the same time: that's how that ring becomes much more stable.
If molecules (in carbonchemistry) have this kind of ring structures with dislocation of electrons, we speak of aromatix molecules.

4. Combinations of different chemical bondings

Until now we spoke about simple ions.
But after treating the main bonds in chemistry (ionic, metalic and covalent), you can understand that there are particles that contain combinations of those bonds.
therefore we look at the so called 'complex ion'.
In complex ions some atoms are connected via covalent bonds.
As a whole they do not make (in this case) a neutral molecule, but an ion with a positive or negative charge.
We know thar in all ions, simple and complex ones, the total number of protons (positive) and the total number of electrons (negarive) are not equal.

examples:

simple:	Na⁺ Cl^- H⁺
complex:	H₃O⁺ OH CO₃^2-

If you know the charge of an ion, you can directly see how the distribution is of protons and electrons in that ion.
example:
An ion with a 2^- charge has two electrons more than protons.
The same for simple and for complex ions.

Question 22
Given the following ions: Ba²⁺ H₃O⁺ Cl^- OH^-
Find in a Periodic Table how many protons and electrons everty ion has.

Complex ions mostly consist of two or more atoms of non metals; they stay toghether within that negative ion through covalent bondingd. Sometimes metal atoms participate in a complex ion)

Question 23
Write the ionic formulas of the five complexions.

Most complex ions contain oxygen + another non metal, and they are connected via covalent bonds.

Question 24
Which suffixes are mostly used for this kind of ions?

If in a substance no metal is present, then you mostly may conclude that this substance has no ions and that the substance is built up of neutral molecules.
All atoms in such a substance are connected via covalent bonds. An exeption is: salts of ammonia.

An example of a substance with different bonding types is calciumcarbonate, CaCO₃.
It is built up of three elements, the metal Calcium and the two non metals Carbon and Oxygen.
The metal Ca will probably be a positive ion, and the rest (CO₃^2-) must be negative.
Between the ions is an ionic bond, but within the negative ion exist covalent bonds between O and C.

Question 25
Give the name and formula of a substance built up of two metals and one non metal, so with two bonding types: the ionic and the covalent bond.

Goto answer 03-25

5. Chemical bond and energy

This subject gives some information of module 7 about chemical reactions, but I believe it is useful and right to ask some attention for it now.

The chemical bond always has to do with valency electrons.
These valency electrons like to connect because they have much energy, and that's how they rearrange to become more stable.
The formation of bonds mostly is exothermic (see module 7).

Below you find a table with bond formation energies.

At each of the bonds, the bond formation energy is given in kJ/mol.
If that is a high value, then the formation of this bond out of the elements seems to deliver lots of energy.

Example:
If you connect to mol of H-atoms, making 1 mol of H-H bonds, you gain 436 kJ of energy.

VI. Bond formation energy

x 100 kJ/mol

H-H	- 4.36	H-F	- 5.63	N≡N	- 9.45	C-F	- 4.4
F-F	- 1.53	H-Cl	- 4.32	C≡C	- 8.3	C-Cl	- 3.3
Cl-Cl	- 2.43	H-Br	- 3.66	C≡N	- 8.9	C-Br	- 2.8
Br-Br	- 1.93	H-I	- 2.99	C=O	- 8.0	C-I	- 2.4
I-I	- 1.51	H-O	- 4.646	C=S	- 2.6	C-O	- 3.5
H-O (alkanol)	- 4.5	N-H	- 3.9	C=C	- 6.1	H-S	- 3.44
C-C	- 3.5	C-H	- 4.1	O=O	- 4.98	P-H	- 3.22

Question 26
Why should a C-C bond have a higher formation energy than a I-I bond?

6. Intermolecular forces

In case of ionic bonds, the ions are in an ionic lattice. In case of a metalic bond, there is a metalic lattice for the participating particles.
How about molecules?

Inside molecules there are intramolecular forces: the covalent bonds.
But, there are also forces between the molecules, called: intermolecular forces.
e.g.: vanderwaals forces, polar forces, hydrogen bridges.

6.1 VanderWaals forces

The Dutch researcher Dr. van der Waals described certain forces between molecules; attraction between molecules, even when these molecules as non polar.
The attraction between non polar molecules is sometimes called: vanderwaals forces.
But in other books, the vanderwaals forces can include polar as well as non polar attraction.
Never mind. If you just understand that there are intermolecular forces.
Particles attract each other because they have mass.
The masses of particles can, when approaching each other - cause a certain polarity in each other (called: induction).
That is responsible for temporarily dipole forces.
Not that we are talking about strong forces. And also they depend on the number of electrons in the particles and how they can be polarised.
It also depends on how big the particles are.
Generally spoken we may say: the vanderwaals forces are related with the molecular mass.

Question 27
The molecules of the Halogenes have low melting points, but have different melting points.
Use the tables and explain those differences.

Goto answer 03-27

Question 28 is an action task:
[this action task must be done in groups. Sometimes you can realise the task in a lab]

The reactions below have been treated before. You must now treat every reaction according to the question list with six questions.

Magnesium	with	Oxigen
Iodine	with	Calcium
Carbon	with	Oxygen
Iron	with	Sulphur
Hydrogen	with	Oxygen

The question list:

Is the position of the elements left or right in the Periodic Table? Do they have large or small atoms?
What electronegativity has every element? Is that high or low?
Note the atomic ray, the vanderwaals ray and the ionic ray of both atoms.
Calculate ΔE.
Give (if possible) the distribution of charges within the product with δ⁺ and δ^-
Which bonding type can be expected in the product, and why?

6.3 Hydrogen bridge

A Hydrogen bridge is another way for molecules to attract each other and to stay toghether.
This way of attraction has directly to do with the existance of dipole molecules, but it is just a bit more than that:
Apart from the presence of dipole forces, there are H-bridges:

A H-atom^δ⁺ is situated in between two δ^- -atoms.

Such a H bridge looks like an ionic bond, but in reality it has kind of covalent character.

δ^- δ⁺ δ^-
O - H - - - O

A H bridge can exist under the next four conditions:

The whole complex must have a linear structure (all three atoms in one line).

the presence of a H^δ⁺ also is a condition.

The δ^- -atoms must be present, one of them mostly being Oxygen (but not necessarily, it could also be Nitrogen)

Those dashes (- - - ) indicat the bridge: one pair of electrons of Oxygen came under the influence of Hydrogen; that's how kind of covalent bond start to exist.

6.4 The molecular lattice

Apart from the ionic and the metalic lattices, there are molecular lattices.
As the word says: this is a lattice built up of molecules.
In such a molecular lattice, the molecules hold each other in a lattice with there mutual attraction forces. These forces can be: vanderwaals, polar or H-bridge.
The vanderwaals forces depend on the mass of the molecule. The bigger and heavier a molecule, the stronger the molecular lattice.

That's how sugar molecules behave in a lattice to form cristals. The forces that hold those molecules toghether are reasonable big, because the sugar moleculas are rather heavy (and also polar).
So there are also polar forces working here.

Water molecules are much lighter, but very strong polar. They remain toghether, not so much because of the vanderwaals forces, but mor because of the polar forces.
The final effect is that they only form a lattice at 0^oC (freezing point)