3. Bonding Flashcards
1
Q
How do ions in ionic compounds achieve stability?
A
Their electron configuration is the same as a noble gas after they have gained/lost an electron
2
Q
What is an ionic bond?
A
The electrostatic force of attraction between the two oppositely charged ions
3
Q
Will ionic bonds exist in isolation?
A
No, they will always form part of a giant ionic lattice
4
Q
When two ions are more charged, what will this mean for the forces of attraction between them?
A
The more charged the ions are, the greater force of attraction between them
5
Q
Are ionic compounds ever molecules?
A
No
6
Q
What does the formula ‘NaCl’ represent?
A
That there is a one to one ratio of sodium to chloride ions
7
Q
What is the formula for an ammonium ion?
A
NH4+
8
Q
What is the formula for a chlorate ion?
A
ClO3-
9
Q
What is the formula for a cyanide ion?
A
CN-
10
Q
What is the formula for a carbonate ion?
A
CO3 2-
11
Q
What is the formula for a chromate ion?
A
CrO4 2-
12
Q
What is the formula for a dichromate ion?
A
Cr2O7 2-
13
Q
What is the formula for a hydrogen carbonate ion?
A
HCO3 -
14
Q
What is the formula for a hydrogen phosphate ion?
A
HPO4 2-
15
Q
What is the formula for a hydrogen sulfate ion?
A
HSO4 -
16
Q
What is the formula for a hydroxide ion?
A
OH-
17
Q
What is the formula for a nitrate ion?
A
NO3 -
18
Q
What is the formula for a nitrite ion?
A
NO2-
19
Q
What is the formula for a manganate ion?
A
MnO4-
20
Q
What is the formula for a peroxide ion?
A
O2 2-
21
Q
What is the formula for a phosphate ion?
A
PO4 3-
22
Q
What is the formula for a sulfate ion?
A
SO4 2-
23
Q
What is the formula for a sulfite ion?
A
SO3 2-
24
Q
What is the formula for a thiosulfate ion?
A
S2O3 2-
25
What is a covalent bond?
A shared pair of electrons, where each electron come from a different atom
26
What happens when a pair of electrons is shared?
A single bond is formed
27
How is a double or triple covalent bond formed?
If two or three pairs of electrons are shared
28
What type of covalent bond does a H2 molecule have?
Single covalent, H-H
29
What type of covalent bond does an O2 molecule have?
Double covalent, O=O
30
What type of covalent bond does an N2 molecule have?
Triple covalent, N≡N
31
What is a coordinate bond?
When both electrons come from the same atom
32
What is a coordinate bond also known as?
A dative covalent bond
33
What are ammonium ions formed from?
NH3 and H+
34
How is an ammonium ion formed?
* NH3 molecule has a lone pair of electrons, and H+ ion has no electrons
* the lone pair in NH3 donates electrons to vacant orbital in the hydrogen (COORDINATE BOND FORMED)
35
How is a coordinate bond represented in displayed formula?
As an arrow, showing where electrons have been donated
36
What are the types of covalent bonding?
* simple molecular
| * giant covalent (macromolecular)
37
Where does metallic bonding exist?
Within a metal, as there is no bonding between metals
38
Why do electrons in metallic bonding become delocalised?
There are no non-metal atoms to transfer electrons to
39
What does the number of delocalised electrons in metallic bonding depend on?
The number of outer electrons the metal atom has
40
What happens in metallic bonding when there are more delocalised electrons?
Increased attracted between ions and electrons as ions more strongly charged
41
What is a metallic bond?
In a metal, when positive metal ions and delocalised electrons attract each other and hold the structure together
42
Do metallic bonds exist in isolation?
No - they form part of a giant metallic lattice
43
What is the strength of ionic, covalent and metallic bonds?
Very strong
44
What are the four types of crystal structure?
* ionic
* metallic
* macromolecular (giant covalent)
* (simple) molecular
45
Why can ionic compounds dissolve in water?
Positive ions in lattice are attracted to opposite charged atom in water and vice versa (as water is a polar molecule)
46
Why can ionic compounds conduct electricity and heat when molten or in aqueous solutions?
Positive and negative ions are free to move and therefore carry charge
47
Why do ionic compounds have high melting points?
Lots of electrostatic forces of attraction between positive and negatively charged ions, which take a lot of energy to break
48
Why are ionic compounds brittle?
If it is bent then layers slide over each other so that like charges are next to each other, then repel and break apart
49
Why don’t simple covalent compounds conduct electricity or heat?
There are no free charged particles to carry charge - only ends of molecules are slightly charged so they are neutral
50
Why do simple covalent compounds have low melting points?
There are weak intermolecular forces between molecules which need little energy to break
51
Why don’t giant covalent compounds conduct electricity? Is there an exception to this?
No charged particles that are free to move
| graphite conducts - 1 free electron per atom
52
Why do giant covalent compounds have high melting points?
Lots of strong covalent bonds which take a lot of energy to overcome
53
Why are giant covalent compounds not soluble in water?
No charged particles - all carbon are neutral so don’t attract to H2O molecules
54
Why do metals conduct electricity and heat?
Have delocalised electrons that can carry a charge
55
Why do metals have high melting points?
Electrostatic forces - ‘metallic bonds’ - need a lot of energy to be broken
56
Why are metals malleable and ductile?
Layers of ions can slide past each other easily - bonds flexible although strong
57
What substances are typical of monatomic structures?
Elements: group 0
58
What substances are typical of simple molecular structures?
* Elements: H₂, O₂, N₂, F₂, Cl₂, Br₂, I₂, At₂, S₈, P₄
| * Compounds: different non-metals bonded
59
What substances are typical of giant covalent structures?
* Elements: Silicon, carbon (diamond and graphite)
| * Compounds: SiO₂
60
What is the structure in monatomic substances?
Individual atoms with very weak forces between them
61
What substances are typical of ionic structures?
Compounds: metal and non-metal bonded
62
What substances are typical of metallic structures?
Elements: one metal
63
What is the structure in simple molecular substances?
Individual molecules with weak forces between them
| atoms within molecules are joined by covalent bonds
64
What is the structure in giant covalent substances?
Lattice structure in which all atoms are joined to other atoms by covalent bonds
65
What is the structure in ionic substances?
Lattice structure of +ve and -vely charged ions
| held together by ionic bonds
66
What is the structure in metallic substances?
Lattice structure of metal ions with outer shell electrons free to move through the structure
67
What are the particles in monatomic substances?
Atoms
68
What are the particles in simple molecular substances?
Molecules
69
What are the particles in giant covalent substances?
Atoms
70
What are the particles in ionic substances?
Ions
71
What are the particles in metallic substances?
Ions and delocalised electrons
72
What is the formula for monatomic substances?
Just the symbol e.g. Ar
| from periodic table
73
What is the formula for simple molecular substances?
Need to learn common examples e.g. H₂O, CO₂, CH₄
74
What is the formula for giant covalent substances?
* Elements: just the symbol
| * Compounds: ratio of atoms e.g. SiO₂, C, Si (need to learn common examples)
75
What is the formula for ionic substances?
Need to learn common ions and work out formula
e.g. MgCl₂ - ratio of Mg:Cl ions is 1:2
76
What is the formula for metallic substances?
Just the symbol e.g. Fe
| from periodic table
77
What is the type of structure and bonding in sodium chloride?
Ionic
78
What are the particles present in sodium chloride?
Ions - Na+ and Cl-
79
Explanation of bonding in sodium chloride?
+vely charged Na ions attracted to -vely charged Cl ions to create ionic bond
80
Explanation for melting/boiling point of sodium chloride?
High - strong ionic bonds take a lot of energy to overcome
81
Explanation for conductivity of sodium chloride?
Conducts when molten or dissolved in water - ions need to be free to move
82
Why is sodium chloride brittle?
When shifted, ions will be next to same charged ion and repel naturally
83
What is the type of structure and bonding in magnesium?
Giant metallic lattice of Mg ions - metallic bonding within the metal
84
What particles are present in magnesium?
Ions and delocalised electrons
85
Explanation of bonding in magnesium?
Attraction between +vely charged ions and -vely charged electrons - electrostatic forces of attraction
86
Explanation for melting/boiling point of magnesium?
High - electrostatic forces of attraction take a lot of energy to overcome
87
Explanation for conductivity of magnesium?
Can conduct as delocalised electrons can move and carry a charge
88
What is the type of structure and bonding in diamond?
Macromolecular - carbon atoms
89
What are the particles present in diamond?
Carbon atoms
90
Explanation of bonding in diamond?
Each carbon atom is bonded to four other carbon atoms covalently
91
Explanation for melting/boiling point of diamond?
High - many strong covalent bonds - lots of energy to break - and no intermolecular forces
92
Explanation for conductivity of diamond?
Doesn't conduct - neutral as not charged, and no free electons - no moving particles
(however good thermal conductor)
93
What is the type of structure and bonding in graphite?
Macromolecular - giant covalent structure
94
What are the particles present in graphite?
Atoms
95
Explanation of bonding in graphite?
* each carbon atom joined to three others covalently
* carbon atoms form layers with a hexagonal arrangement (layers have weak forces in between)
* each atom has one non-bonded, delocalised electron
96
Explanation for melting/boiling point of graphite?
High - covalent bonds
97
Explanation for conductivity of graphite?
Conducts - delocalised electron per atom is free to move
98
What is the type of structure and bonding in iodine?
Covalent, simple molecular, with weak van de waal forces in between
99
What are the particles present in iodine?
Two iodine atoms - diatomic iodine molecules
100
Explanation of bonding in iodine?
Iodine molecules joined by weak van de waal's (intermolecular) forces
101
Explanation for melting/boiling point of iodine?
* m.p. - 113.7 and b.p. - 183.35
| * low - van de waals forces are weak so don't need as much energy to break
102
Explanation for conductivity of iodine?
Doesn't conduct - covalently bonded so no delocalised electrons
103
Can iodine undergo sublimation?
Yes
104
What is the type of structure and bonding in ice?
* Covalent (simple molecular)
| * With intermolecular force - hydrogen bond
105
What are the particles present in ice?
H₂O molecules - hydrogen and oxygen covalently bonded
106
Explanation of bonding in ice?
Slightly +ve regions of one molecule attracted to slightly -ve regions of another to form hydrogen and vice versa
107
Explanation for melting/boiling point of ice?
* b.p. - 100 and m.p. - 0
| * relatively high for a simple molecular substance due to hydrogen bonds
108
Explanation for conductivity of ice?
* using distilled water doesn't conduct - no charged particles that are free to move
* tap water contains ions
109
Why is ice less dense than water?
Molecules are arranged further away in ice
110
What is the shape of a simple molecule and ion determined by?
The number and type of electron pairs that surround the central atom
111
How do charge clouds try and arrange themselves?
In a way that minimises the repulsion between charge clouds - by moving as far apart as possible
112
What is the shape name and therefore bond angle of CO₂?
Linear - 180°
113
Why is CO₂ a linear molecule?
It has double bonds
114
What are the two types of electron pairs in a molecule?
* bonding pairs
| * non bonding / lone pairs
115
What are bonding pairs of electrons?
Those that are involved in a chemical bond
116
What are non-bonding/lone pairs of electrons?
Those that are not involved in a chemical bond
117
Which type of electron pair have a greater repulsion in a molecule?
Lone pairs
118
How does repulsion increase, in terms of electron pairs?
bonding- bonding pair < lone - bonding pair < lone - lone pair
119
What are the steps taken to predict the shape of a molecule?
1. determine no. of electron pairs
2. identify number of lone pairs
3. use these to determine shape of molecule and bond angle
120
How do you calculate the number of electron pairs surrounding the central atom?
(number of outer electrons of central atom (group no.) + no. of atoms bonded) ÷ 2
121
How do you calculate the number of lone pairs?
Number of electron pairs - number of atoms bonded
122
How much do lone pairs reduce bond angle by?
Approximately 2.5°
123
Why do lone pairs reduce bond angles?
Due to increased repulsion
124
Work out the shape and bond angle of CH₄.
Tetrahedral, 109.5°
125
Work out the shape and bond angle of H₂O.
V-shaped, 104.5°
126
What is the shape and bond angle of a molecule with two bonding, and no non-bonding electron pairs?
Linear, 180°
127
What is the shape and bond angle of a molecule with three bonding, and no non-bonding electron pairs?
Trigonal planar, 120°
128
What is the shape and bond angle of a molecule with two bonding, and one non-bonding electron pair?
Bent (v-shaped), 117.5°
129
What is the shape and bond angle of a molecule with four bonding, and no non-bonding electron pairs?
Tetrahedral, 109.5°
130
What is the shape and bond angle of a molecule with three bonding, and one non-bonding electron pair?
Trigonal pyramidal, 107°
131
What is the shape and bond angle of a molecule with two bonding, and two non-bonding electron pairs?
Bent (v-shaped), 104.5°
132
What is the shape and bond angle of a molecule with five bonding, and no non-bonding electron pairs?
Trigonal bipyramidal, 120° and 90°
133
What is the shape and bond angle of a molecule with four bonding, and one non-bonding electron pair?
Trigonal pyramidal or see-saw, 120° and 90°
134
What is the shape and bond angle of a molecule with three bonding, and two non-bonding electron pairs?
Trigonal planar or t-shape, 120° and 90°
135
What is the shape and bond angle of a molecule with two bonding, and three non-bonding electron pairs?
Linear, 180°
136
What is the shape and bond angle of a molecule with six bonding, and no non-bonding electron pairs?
Octahedral, 90°
137
What is the shape and bond angle of a molecule with five bonding, and one non-bonding electron pair?
Square based pyramid, 90°
138
What is the shape and bond angle of a molecule with four bonding, and two non-bonding electron pairs?
Square planar, 90°
139
What is the shape and bond angle of a molecule with three bonding, and three non-bonding electron pairs?
T-shaped, 90°
140
What is the shape and bond angle of a molecule with two bonding, and four non-bonding electron pairs?
Linear, 180°
141
What is affected when determining molecule shape in the case of an ion?
The number of electrons of the central atom is affected by its charge
142
When determining the shape of a positive ion, what has to be done to the number of outer electrons of the central atom?
Subtract number of electrons lost from number of outer electrons
143
When determining the shape of a negative ion, what has to be done to the number of outer electrons of the central atom?
Add number of electrons gained to number of outer electrons
144
Work out the shape and bond angle of an NH4+ ion.
Tetrahedral, 109.5°
145
Work out the shape and bond angle of a CIF4- ion.
Square planar, 90°
146
What is affected when working out the molecule shape of an ion instead of an atom?
The number of electrons of the central atom is affected by its charge
147
How can the molecule shape of a positive ion be worked out?
* the central atom has lost that no. of electrons
| * so the charge needs to be subtracted from the no. of outer electrons of the central atom
148
How can the molecule shape of a negative ion be worked out?
* the central atom has gained that no. of electrons
| * so the charge needs to be added to the no. of outer electrons of the central atom
149
Why is the bond angle in an amide ion smaller than that in an ammonia molecule?
* amide ion has 2 lone pairs; ammonia has one
| * more repulsion between two lone pairs and pushes them closer together
150
What is electronegativity?
The power of an atom to attract the pair of electrons in a covalent bond
151
What does electron density describe?
How the negative charge in an atom is distributed
152
What does electronegativity depend on?
* nuclear charge (no, of protons)
* atomic radius
* shielding
153
What is atomic radius?
The distance between the nucleus and outer shell electrons
154
What happens to electronegativity as nuclear charge increases?
It increases
155
What happens to electronegativity as atomic radius increases? Why?
It decreases - bonding pair of electrons is further from the nucleus so less attracted to charge of the nucleus
156
What happens to electronegativity as shielding increases?
It decreases
157
What happens to electronegativity across a period?
It increases
158
Why does electronegativity increase across a period?
* proton no. ↑ and size ↑ so nuclear charge ↑
| * shielding doesn't ↑ as electron shells are the same
159
What happens to electronegativity down a group?
It decreases
160
Why does electronegativity decrease down a group?
* shielding ↑ (as no. of shells ↑)
| * and distance from nucleus (atomic radius) ↑
161
What is the most electronegative element?
Fluorine
162
Is there a difference in electronegativity in a covalent bond with two atoms that are the same?
No
163
Why is electron density evenly distributed in a covalent bond between two atoms that are the same?
Because there is no difference in electronegativity
164
What results in a non-polar bond?
When there is an equal sharing of electrons due to having no difference in electronegativity
165
What happens in a covalent bond between two atoms that are the same, in terms of electronegativity?
A non-polar bond is formed due to no difference in electronegativity
166
What happens in a covalent bond between two atoms that are the different, in terms of electronegativity?
A polar bond is formed, due to a difference in electronegativity and therefore the shared electrons not being evenly distributed
167
What does the difference in electronegativity between atoms in a covalent bond result in?
The shared electrons not being evenly distributed, and therefore a polar bond
168
When there is unequal sharing of electrons in a covalent bond, which atom will the electron pair be drawn towards?
The atom that is more electronegative
169
When does a polar bond form?
* In a covalent bond between two atoms that are different
* difference in electronegativity therefore unequal sharing of electrons
* electron pair drawn towards more electronegative atom = partial charges
170
What does a larger difference in electronegativity mean for a polar bond?
The bond is more polar
171
How is a polar bond thought to be, in terms of character?
Covalent with some ionic character
172
Can ionic bonds show covalent character?
Yes
173
How can ionic bonds show covalent character?
The ionic nature of the bond is reduced as the electron cloud is distorted by strong charges of the ion
174
What happens if a cation has a high positive charge or small size?
It will tend to attract electrons towards itself
175
What happens if an anion has a high negative charge or large size?
It will have an electron cloud that is easily distorted
176
What will it mean for an ionic bond if ions have high charges?
The polarity of the bond increases, giving the bond covalent character
177
What does the degree of ionic or covalent character depend on?
The atoms involved
178
What would happen if there were no forces holding water molecules together?
They would move apart and become a gas
179
What are the forces that exist between covalent molecules called?
Intermolecular forces
180
How strong are intermolecular forces compared to covalent bond?
Between 1/10 and 1/100 of the strength
181
What are the three types of intermolecular forces?
1. Van der Waal's
2. permanent dipole-dipole forces
3. hydrogen bonds
182
Where do Van der Waal's forces exist?
Between all molecules where they're liquid or solid
183
Why is there an imbalance in charge on a molecule, even though there may be no difference in electronegativity?
At any instant the electron distribution in a non-polar covalent bond can be asymmetrical due to constant movement of electrons
184
What is a temporary dipole?
When there is an imbalance of charge on a molecule due to the electron distribution in a non-polar bond being asymmetrical because of constant electron movement
185
What does a temporary dipole do to the adjacent molecule?
Induces an opposite dipole - and this continues through the structure
(this is known as the van der waal's force)
186
How are van der waal's forces formed?
When there is an asymmetrical distribution of electrons in a molecule which leads to an imbalance in charge - and this induces an opposite charge on the next molecule
187
What are van der waal's forces also known as?
Temporary dipole-dipole forces
188
Can van der waal's form in monoatomic substances?
Yes
189
What is the only intermolecular force in diatomic and monoatomic molecules?
Van der waal's - but they occur between all molecules
190
Are van der waal's forces in addition to any other intermolecular force?
Yes
191
What is the strength of van der waal's forces dependant on?
* no. of electrons present/size of molecule
| * shape of the molecule
192
What does a molecule having more electrons mean for the strength of its van der waal's forces?
more electrons = stronger van der waal's
193
Why does having more electrons increase the strength of van der waal's forces?
* the outer electrons are further from the nucleus and so attracted less strongly by the nucleus
* so temporary dipoles are easier to induce
194
What does a branching mean for the strength of its van der waal's forces?
When molecules have branches the van der waal's forces are weaker
195
Why does branching decrease the strength of van der waal's forces?
The molecules are further away when branched, so straight chain alkanes pack closer together
196
What is the weakest type of attraction that can exist between molecules?
Van der Waal's
197
What are polar molecules?
Molecules where there is a difference in electronegativity between the atoms involved
198
Which type of molecules have permanent dipoles?
Polar molecules
199
What type of dipole do polar molecules have?
Permanent
200
How do polar molecules attract each other?
Each molecules has opposite charges on each end; these opposite charges on the different molecules attract eachother
201
Are permanent dipoles induced?
No, they already exist
202
Why are permanent dipoles not constantly changing?
Because their dipoles already exist due to differences in electronegativity
203
How will molecules with permanent dipoles arrange themselves?
So that oppositely charged ends are closest
204
Do all polar molecules have polarity overall?
No
205
Which polar molecules do not have overall polarity?
Symmetrical molecules
206
Why do symmetrical molecules not have polarity overall?
The dipoles can cancel each other out to leave the molecule with no overall polarity
207
What is meant by a symmetrical molecule?
A molecule where the dipoles are the same around a central atom or bond
208
What is the strongest intermolecular force that can exist between symmetrical polar molecules?
Van der Waal's
209
Which are stronger: Van der Waal's or permanent dipole-dipole?
Permanent dipole-dipole
210
What is the strength of a permanent dipole dependant on?
The difference in electronegativity of the atoms in the polar bond
211
What will happen to a permanent dipole-dipole attraction when there is a larger difference in electronegativity?
There will be greater partial charges and therefore a stronger permanent dipole-dipole attraction
212
Which atoms have a large difference in electronegativity to hydrogen atoms?
Nitrogen, oxygen and fluorine
213
What are the three most electronegative atoms?
Nitrogen, oxygen and fluorine
214
What does the large difference in electronegativity between hydrogen and either nitrogen, oxygen or fluorine mean?
Very polar bonds are formed, and therefore a strong permanent dipole
215
What is the strongest type of permanent dipole?
Hydrogen bonding
216
What is a stronger; a hydrogen bond or a covalent bond?
Covalent
217
What is the least electronegative atom?
Hydrogen
218
When are strong permanent dipoles formed?
When there is a large difference in electronegativity between hydrogen and either fluorine, oxygen or nitrogen
219
How is a hydrogen bond different to a dipole-dipole force?
The nitrogen, oxygen or fluorine atom can draw hydrogen's single electron towards itself
220
What is the result of nitrogen, oxygen or fluorine drawing hydrogen's single electron towards them?
Hydrogen exposes its single proton and therefore a strong δ+ hydrogen end and a strong δ- end
221
What is required for a hydrogen bond to form?
A lone pair of electrons on the N, O or F
222
Why is it important that the non-hydrogen atom has a lone pair to form a hydrogen bond?
So this lone pair can strongly attract the δ+ end of hydrogen - which hydrogen can't repel without any non-bonding electrons
223
Why can't hydrogen repel the non-hydrogen atom in a hydrogen bond?
It has no non-bonding electrons
224
What is a hydrogen bond?
The attraction between the strong δ+ hydrogen end and the lone pair of electrons on a N, O or F atom
225
What does a hydrogen bond have characteristics of?
A dipole-dipole force as well as a covalent bond
226
How strong is a hydrogen bond in comparison to a covalent bond?
1/10 the strength of a covalent bond
227
What is the strongest type of intermolecular force?
Hydrogen bonds
228
What are the three intermolecular forces responsible for keeping the structure of hair in place?
* hydrogen bonds
* salt linkages
* disulfide bonds
229
How do hydrogen bonds affect the structure of molecules in hair?
Hydrogen bonds attract the long keratin molecules to each other and them into a spiral shape
230
How can the corresponding strength of intermolecular forces be represented as?
Van der waal's < permanent dipole-dipole < hydrogen
231
When will a substance turn from a solid to a liquid?
When there is enough heat energy to break enough intermolecular forces
232
What is melting temperature?
The heat energy required to break enough intermolecular forces to turn a substance into a lquid
233
What is boiling temperature?
The energy required to break enough intermolecular forces to turn that substance into a gas
234
What do substances with hydrogen bonds tend to have higher boiling points than?
Substances with permanent dipole-dipole bonds, and those with van der waal's
235
What is the pattern of boiling points of the hydrides in periods 2-5?
Apart from the 1st in the period, there is a general increase across the period
236
Why does the boiling point of hydrides in groups 2-5 increase across the period?
* molecules are increasing in size
* due to more electrons
* greater polarization and dipole increases in strength
* (van der waal's also stronger due to bigger molecule)
237
In a graph showing boiling point of hydrides between periods 2-5, how and why do H2O, HF and NH3 not follow the trend?
* how? they are much higher than the other hydrides in their group
* why? there are hydrogen bonds between molecules
238
Why does H2O have a greater boiling point than HF and NH3?
* H2O - equal no. of hydrogens and oxygens so more bonds - also 2 lone pairs on oxygen so quite -ve attraction between H&O
* HF - strong bonds (F is most electronegative) but unequal ratio of H to lone pairs
* NH3 - lower b.p. as less electronegative & unequal ratio of N to H
239
Why are the boiling points of group 4 hydrides lower than those in 5,6, and 7?
Molecules in group 4 have symmetrical tetrahedral shape - so electronegativity cancels out = no permanent dipole - only have van der waal's
240
What happens to molecules when water is heated?
They gain energy, moving faster and further apart as a result of hydrogen bonds being broken
241
What happens to water's volume and density as it is heated?
Water becomes less dense and volume increases
242
What happens to water's volume and density as it is cooled?
Water becomes more dense UNTIL it reaches 4°C
243
When is water at its most dense?
4°C
244
Why does water float and expand when it is frozen?
Hydrogen bonds are quite long and therefore hold the water molecules slightly further apart
