2 - Atoms, Reactions, Bonding And Structure Flashcards
1
Q
Define binary compound
A
A compound containing two elements only
2
Q
Define Polyatomic ion
A
An ion containing more than one atom
3
Q
What happens when a ionic substance dissolves in water
A
Positive and negative ions separate and become hydrated.
4
Q
In reactions involving ionic compounds dissolving in water some ions may not be involved.
What are they called
A
Spectator ions
5
Q
Acid + hydroxide
Ionic equation
A
H+ (aq) + OH- (aq). —> H2O(l)
6
Q
Acid carbonate (ionic equation)
A
2 H+ (aq) + CO3 2- (aq). —> H20 (l) + CO2 (g)
7
Q
Acid + hydrogen carbonate
(Ionic equation)
A
H+ (aq) + HCO3 - (aq) —> H2O(l) + CO2 (g)
8
Q
Acid + ammonia
(Ionic equation)
A
H+ (aq) + NH3 (aq) —> NH4+ (aq)
9
Q
Substance + oxygen -?
A
Oxides
10
Q
Metal + water — ?
A
Metal hydroxide + hydrogen
11
Q
Metal + acid —
A
Salt + hydrogen
12
Q
Oxide +acid —
A
Salt + water
13
Q
Hydroxide + acid —?
A
Salt + water
14
Q
Carbonate + acid —?
A
Salt + water + carbon dioxide
15
Q
Ammonia + acid —?
A
Ammonium salt
16
Q
Metal carbonate on heating —?
A
Metal oxide + carbon dioxide
17
Q
What is avogadros constant
A
6.02 x 10^23 mol-1
18
Q
Why was avogadros number chosen
A
So that the mass of one mole of particles of a substance equals the mr in grams.
19
Q
Define isotopes
A
– Atoms of the same element with different numbers of neutrons and different masses. Isotopes have the same atomic number but a different mass number
20
Q
Isotopes always have- — - - - - chemical reactions as they have - - - - - electron configuration but physical properties like density may be — - - - -
A
The same
The same
Different
21
Q
What can make an isotope radioactive
A
Sometimes the extra neutrons make the nucleus big and unstable
22
Q
Isotopes of an element always have the…..
A
same atomic number but a different mass number due to having extra neutrons in the nuclei.
23
Q
Relative atomic mass
A
The weighted mean mass of an atom of an element relative to one twelfth of the mass of an atom of carbon-12
24
Q
Relative isotopic mass
A
the mass of an isotope relative to one twelfth of the mass of an atom of carbon-12
25
Relative molecular mass
compares the mass of a molecule with the mass of an atom of carbon-12
26
How do u calculate mr
Adding the Ar of all the elements
27
How do you calculate the relative formula mass
Adding Ar of all the elements in the empirical formula
28
Define relative formula lass
Compared the mass of a formula unit with the mass of an ato, of carbon 12
29
What is relative atomic mass calculated from
Using the isotopic masses and their abundence
30
How are isotope abundances found
Using a mass spectrometer
31
What is in the x axes of a mass spectrometer
m / z
32
What is in the y axes of a mass spectrometer
Percentage abundance
33
How to calculate relative atomic mass
(Atomic mass 1 x% abundance) + (atomic mass 2 x % abundance )
———————————————————
100
34
How to use a mass spectrometer
1- sample In
2- vaporised then ionised to form positive ions
3- ions accelerated - heavier move slower
4- ions detected on mass spectrum as mass: charge ration m/z - as they reach detector they create a signal - detected so the greater the abundance the larger the signal
35
How to find relative atomic mass without knowing the % abundance (isotopes)
(Mass) (% - written as x) + (mass2) (% - written as 100-x)
———————————————— — —
100
1) expanse + simplify
2- work out x
3 - work out 100 - x
36
Orbitals definition
A region of space where there is a high probability of finding an electron
37
How many electrons can orbitals have
Up to 2 that spin in opposite directions
38
What does isoelectronic mean
Same electronic configuration
39
Electrons charge
-1
40
Where are electrons located
In orbitals
These orbitals take up most the volume of the atom
41
Where is most the mass of the atom concentrated
In the nucleus
42
What subatomic particles are located in the nucleus
Protons and neutrons
43
Relative mass and charge of proton
Rm = 1
Charge = +1
44
Relative mass and charge of neutron
Rm= 1
Charge = 0
45
Relative mass of an electron
1
——
1836 (check)
46
How do you find number of neutrons
Mass number minus atomic number
47
Do negative ions have more electrons or protons
Electrons
48
What decides the chemical properties of an element?
So how does this effect isotopes?
The number and arrangement of eelctromsn
They have the same chemical properties as they still have the same electron configuration
49
What did some ancient Greeks use to think about atoms?
All matter was made from indivisible particles
50
At the start of the 19th century how did John dalton describe atoms
Sold spheres
And different types of spheres make up differnt elements
51
What did JJ Thomson prove + how
Atoms weren’t sold and indivisible.
His measurements of charge and mass showed that an atom must contain even smaller , negatively charged particles (he called them ‘corpuscles’ - we call them electrons)
The new model was known as the plumb pudding model - a positively charged sphere with negativley charged electrons embedded in
52
When did JJ Thomson do his experiments
1897
53
When did John dalton describe atoms as solid spheres
19th century
54
When did Ernest Rutherford do the gold foil experiment and with who?
1909
His students - hans Geiger and Ernest marsden
55
What was the gold foil experiment
And what were they expecting to happen becuase of the plumb pudding model
They fired alpha particles (which are Moseley postives charge) as an extremely thin sheet of gold.
Expected most of the alpha particles to be deflected very slightly by the positive pudding that made up most of the atom
56
What actually happened in the gold foil experiment
Most the alpha particles passed straight through the gold atoms and a very small number where deflected backwards - proved the plum pudding model was wrong
57
What did Rutherford replace the plumb pudding model with
The nuclear model
58
What was Rutherfords nuclear atom
1) a tiny positively charged nucleus at the centre of the atom, where most of the atoms mass is concentrated
2) the nucleus is surrounded by a cloud of negative electrons
3) most of the atom is empty space
59
What did Henry Moseley discover
The charge of the nucleus increased from one element to another in units of one.
Which led to him discovering it contained positively charged particles called protons.
60
Why did James Chadwick continue to experiment and What did James Chadwick discover
There was a problem with the previous model - the nuclei of the atoms was heavier than it would be a if it just contained protons.
James predicted that there was another subatomic particle in the nucleus wuth mass but no charge - he discovered the neutron
61
Why did Neil bohr continue to experiment
Scientists realised that electrons in a cloud around the nucleus of an atom would spiral down into the nucleus causing the atom to collapse
62
What 4 basic principles where in Niel bonds new model of the atom
1) electrons can only exist in fixed orbits or shells, and not anywhere in between
2) each shell has a fixed npenergy
3) when an electron moves between shells electromagnetic radiation is emitted or absorbed
4) becasue the energy of shells is fixed, the radiation will have a fixed frequency
63
How did Bohr mod well explain why some elements are inert (noble gasses)
The shells of an atom can only hold fixed number of electron,and an elements reactivity is due to its electrons.
When an atom has full shells of electrons it is stable and doesn’t react
64
What does the most accurate model of the atom we know today involve
Inc positive and negatives
Complicated quantum mechanics.
You can never know where an electron is or which direction it is goin inn at any moment, but you can say how likely it is to be at any particular point in the atom,m
P- more accurate and explains some observations that Bohr model can’t
N-more complicated visually
65
When is relative molecular mass used
When referring to simple molecules
66
How do you find relative molecular mass
Add up the relative atomic mass values of all the atom is in the molecule
67
When is relative formula mass used
For compounds that are ionic (or covalent)
68
How do you find the relative formula mass
Add up the relative atomic masses of all the ions in the formula unit
69
What is a mass spectra produced by
A mass spectrometer
70
What is a mass spectrometer
A device which are used to find out what samples are made up of by measuring the masses of their components
71
What can a mass spectra tell us
The relative isotopic masses and abundances of differnt elements
72
What is on the y-axis for a mass spectra
The abundance of ions as a %
- for an element the height of each peak gives the relative isotopic abundance
73
What is on the x-axis for a mass spectra
Units given as ‘mass/charge’ ration
(m /Z)
Since the charge on most ions is +1 you can often assume this axis is the relative isotopic mass
74
What can the mass spectra be used to work out
The relative atomic masses of differnt element
75
How to calculate relative atomic masses using a mass spectra
1) multiply each relative isotopic mass by its relative isotopic abundance and add up the results
2) divide by the sum of the isotopic abundances
76
What is avogadro constant
6.022 x 10 23^
77
Formula for finding number of moles from the number of atoms / molecules
Nu. Moles= nu. of particles you have / nu.Particles in a moles
78
T or f
Is molar mass the same as relative molecular mass
T
79
What is molar mass units
G mol-1
80
Moles formula
Mass /mr
81
What is the order of the orbitals
1s2 2s2 2p6 3s2 3p6 4s2
82
How many electrons in energy level 1
2
83
How many electrons in energy level 2
8
84
How many electrons in energy level 3
18
85
How many electrons in energy level 4
32
86
What are the differnt types of orbitals
S
P
D
F
87
What shape is orbital s
Spherical
88
What shape is orbital p
Dumb-bell shell
89
What are the rules of filling orbitals with electornx
Aufbau principle
Fill in order of inc energy
90
What is paulis exclusion principle
Within an orbital, electrons pair up with opposite spin so that the atom is as stable as possible.
Electrons in the same orbital must have opposite spin
91
What are the 3 rules for writing out electron configuration
1) lowest energy orbitals filled first
2) one electron occupies each orbital before pairing begins (hands rule)
3) no single orbital holds more than 2 electrons
92
When doing electron configuration for ions, they will lose electron in reverse order except what
4s will empty before 3D
93
If electrons are unpaired and therfore unbalanced it produces what
Natural repulsion between the electrons making the atom very unstable
The electrons may take on differnt a range,ents to improve stability
94
What is ionic bonding
Electrostatic forces of attraction between oppositely charged ions
95
What 2 elements are exceptions to the rule about electrons filling up orbitals
Copper and chromium
As half filled sub shells are more stable
96
How are water molecules part of the crystalline structure represented by
XXX.H20
97
What are the units for volume
1dm 3
98
Dm3 = xxxxxx cm3
1000
99
Concentration and volume equation
Conc. = mole / vol
100
What are the units for concentration
Mol / dm3
M
101
What does standard solution mean
+how
A solution of a known concentration.
An exact mass of solute is dissolved in a solvent and made up to an exact volume of solution
102
How do you go from mol/dm3 to g/dm3
Mol/dm3 x molar mass
103
What is avogadros law
At the same temperature and pressure, equal volumes of differnt gases contain the same number of molecules
104
What is molar volume
The molar gas volume Vm is the volume per mole of gas molecules at a stated temperature and pressure
105
At room temperature the molar gas volume is
24.0 dm3mol-1
106
Gas volume equation
Moles = vol / molar volume(24)
107
What is the method to carry out an experiment to determine the water crystallisation in hydrated crystals
1) weight empty crucible
2) add hydrated salts into weight crucible. Weigh the crucible containing the hydrated salt.
2) using a pipe-clay triangle, surpport the crucible contains the hydrated slat on a tripod. Heat the crucible and contents gently for around a Minuit. Then hear it strongly for another 3 mins
4) leave crucible To cool. Weigh crucible and anhydrous salt.
Calculate
1) calculate amount, in mol, of Anhydrous salt.
2) calculate mass + amount of water
3) find smallest whole number ratio
108
What are the bonds like in ionic bonding
And how does this effect melting and boiling points
Very strong
Attractions between the opposite ions bind the together.
It takes a lot of energy to break these bonds in order for the ions to move apart and change state.
- so the melting and boiling points of such compounds are high (over 500°c)
- the melting and boiling points usually increases as the change of the ions increase
109
Do solid ionic lattice conduct electricty
No
They cannot move as they are held tightly
110
Do liquid / aqueous solutions of ionic lattices conduct electricty
Ions are free to move and so can carry a charge
111
Why are ionic compounds hard
The strong attractive forces holding the lattice together
112
Why are ionic crystals brittle
When hit in the correct place ions are lined up alternatively to attract.
A force can push the rows so same ions now repel and the crystal will flake away
113
Why does it mean that water is a polar solvent
The distribution of electrons in the water is uneven giving it slight positive and slight negative chagre at each end
These charges attract the positive and negative ions and surround them according to charge . The ions are now hydrated , and the compound is dissolved
114
In order for ionic lattice to dissolve what needs to happen
Energy must be provided to break down the lattice . This is provided by the hydrates ion.
The water molecule must then attarct and surround the ions
115
In ionic compounds where ions have larger charges the ionic attraction may be too strong for what?
Water to break down the lattice structure
116
What does solubility of ionic lattice depend on
The relative strength of the attraction within the giant ionic lattice and the attraction between ions and water molecules
117
Where do we see stronger electrostatic forces of attraction between (ionic lattice)
ions with greater charge densities. Charge density is the charge an ion carries comapred to its size,
118
What determines the boiling point of a lattice
The attraction between the positive and negative ion
119
What holds a ionic lattice together
The electrostatic attraction between the positive ions and negative ions and the negative ions produce a strong giant ionic lattice
120
What are ionic compounds generally
Crystals with straight edges.
Suggesting that the ions within the crystal line up alternatively in repeating straight lines
121
What is standard solution
A solution of a known conc.
122
What is avogadros law
At the same temp. And pressure, equal volumes of different gases contain the same number of molecules
123
What is molar volume
The molar gas volume Vm is the volume per mole of gas molecules at a stated temp and pressure
124
What is Dative covalent bonding
Both bonding electrons originate from the same atom
125
What is molar gas volume
The space that one mole of a gas occupies at a certain temp and pressure
126
Units of molar gas volume
Dm3 mol-1
127
At room temp and pressure molar gas volume is 24
Add more info (numbers)
298K (25°c) and 101.3KPa
128
Number of moles =
(Molar gas volume)
Volume in dm3
———————
Molar gas volume
129
What is rhe ideal gas equation
pV = nRT
130
What does the ideal gas equation let you find
The number of moles in a certain volume at any pressure and temperature
131
What do the letters stand for in pV = nRT
+ units
P = pressure (Pa)
v = volume (m ^3)
n = number of moles
R = 8.314 J K-1 mol -1
T = temperature (K) (K = C° + 273)
132
What do balanced equations have
The same number of atoms on both sides
133
How do you know if an ionic equation is wrong
if the charges arnt balanced as well as the atoms
If the charges don’t balance the equation isn’t right
134
What does the reaction stoichiometry tell you
The ratios reactants to products
135
There are lots of ions that are made up of a group of atoms with an overall change
What are these called
Molecular ions
136
Charges in ionic compounds are always …
Balanced
137
What is formed when acids and bases react
Water and salt
138
What are salts
Ionic compounds
139
All solid salts consist of..
In some salts..
a lattice of positive and negative ions.
Water molecules are incorporated in the lattice too
140
What is the water in a lattice called
Water of crystallisation
141
As solid slat contains water of crystallisation is …
Hydrated
142
When is a salt anhydrous
When it doesn’t contain water of crystallisation
143
T or f
Ome mole of a particular hydrated salt always has the same number of moles of water of crystallisation
T
144
What do many hydrated slats loose when heated
Their water of crystallisation
145
The electrons in energy levels are all given numbers known as what?
Primcipal quantum numbers
146
Shells further away from the nucleus have … energy ( and a … principle quantum number)
Higher
Larger
147
What are energy levels further divided into
Sub-shells
148
What are the subshells called
S
P
D
F
149
How many orbitals in s subshell
1
150
How many orbitals in the p subshell
3
151
How many orbitals in the d subshell
5
152
How many orbitals in the d subshell
5
153
How many orbitals in the f subshell
7
154
Max electroms in s subshell
2
155
Max elevtrons in p subshell
6
156
Max electrons in d subshell
10
157
Max electrons in f subshell
14
158
Max electrons in first energy level
2
159
Max electrons in second energy level
8
160
Max electrons in 3rd energy level
18
161
Max electrons in 4th energy level
32
162
T or f
Orbitals with in the same subshell have the same energy
T
163
If there are two electrons in an orbital they have to - —-
+ what is this called
Spin in opposite directions
Spin- pairing
164
What shape are s orbitals
Spherical
165
What shape are p orbitals
Dumbbell shape
166
There are 3 p orbitals and they are at ……. To each other
Right angles
167
Rules for filling up energy levels
1) electrons fill up the lowest energy sub shells first
2) electrons fill orbitals with the same energy singly before they start pairing up
168
Definition of ionincbonding
An ionic bond is an electrostatic attraction between two oppositely charged ions
169
What do dot and cross diagrams show
The arrangement of electrons in an atom or ion.
Each electron is represented by a dot or cross.
They can show which atom the electrons in a bind originally came from
170
Why is a giant ionic lattice called giant
It’s made up of the same basic repeating unit over and over again
171
Why does giant ionic bonding form
Ion is electrostically attracted in all directions to ions of the opposite charge
172
Why do atoms form covalent bonds
To gain a full outer shell
173
What is a covalent bond
A covalent bond is the strong electrostatic attraction. Between a shared pair of electron and the nuclei of the bonded atoms
174
Do all atoms in covalent bonds have 8 electrons in their outer shells
No
Some have less
A few use d orbitals to ‘expand the octet’ - and have more than 8
175
What is average bond enthalpy
Measures energy required to break a specific bond
176
The stronger the covalent bond the …. Energy required to break it so the …. Value of bond enthalpy
More
Greater
177
In covalent bonds
Is there ever more than one pair of electrons
Yes
Some have double or triple
178
What is the ideal gas constant
8.314
179
How to turn c° in K
C° + 273
180
How to turn cm3 into m 3
Divide 1 000 000
181
What can the ideal gas equation be drivied from
Coming bottles law, Charles law + avogadros law
Also can be from kinetic theory of gases
182
What two key assumptions is the ideal gas equation based on
- no intermolecular forces between gas particles
- the volume occupies by the molecules themselves is negligible compared to he volume of the container
183
What other assumptions are made in the ideal gas equation
- gases are made of molecules which are in constant random motion in straight lines
- the molecules behave as rigid spheres
- pressure is sure to collisions between molecules are the walls of the container
- all collisions are perfectly elastic (no loss of kinetic energy )
- the temperature of the gas is proportional to he average kinetic energy of the molecules l
184
What does VSEPR stand for
Valence shell electron pair repulsion theory
185
What is electron pair repulsion theory
Electrons will take up positions as far away from each other as possible to minimise repulsion
186
What can co2 be shape be described as
Linear
187
Explain a linear molecule
2 atoms around a central atom
2 regions of bonded pairs on opposite sides
180°
188
What is the molecule shape when there are 3 atoms around a central atom and with what angle
Trigonal planar shape
189
What are the angles in a trigonal planar shape
120°
190
What does a dotted line when what draw rings molecule shapes
Away
191
What does the wedge mean when drawring the 3D structures of molecules
Towards
192
What is rhe shape when 4 atoms are bonded around the central atom
TetrHedral shape
193
What are the angles in a tetrahedral shape
109.5
194
What is the name of the shape for 5 atoms around the centre atom
Trigonal bipyramidal
195
What are the angles in the Trigonal bipyramidal shape
120°
And
90°
196
What is the name of the shape where there are 6 atoms around the central atom
Octrahedral
197
What are the angles in an octrahedral
All 90°
198
What are the order of biggest to smallest angles in shapes of molecules
Lone pair / lone pair angles are the biggest
Lone pair / bonding pair angles are the second biggest
Binding pair / bonding pair bond angles are the smallest
199
What is the way of predicting molecular shape known as
Electron pair repulsion theory
200
What are the bond angles in methane and why
No lone pairs so all angles 109.5°
201
What are the angles of ammonia and why
1 lone pair
All 3 angles are 107°
202
Water angles in molecular shape and why
2 lone pairs reduce the bond angle
104.5°
203
How do you predict the shape of a molecule
Find the central atom work out the number of electrons in the outer shell of the central atom.
Then work out how many electrons are shared with the central atom
Add up the electrons and divide by 2 to find the number of electron pairs on the central atom and if you have an ion remember to count for its charge
compare the number of electron pairs with a number of bonds to find the number of lone pairs.
You can then use the number of electron pairs and the number of lone pairs and bonding centres around the central atom to work at the shape.
204
What is rhe shape when the era re no lone pairs but 4 electron pairs around the centrato,
Tetrahedral
205
What is the shape when there is one lone pair and 4 electron pairs around th e central atom
Trigonal pyramidal
206
What is the shape when there are 2 lone pairs and 4 electron pairs around the central atom
Nonlinear or bent
207
What is theoretical yield
Mass of products that should be made in a reaction if no chemicals are ‘lost’ in the process.
208
What is percentage yield
The actually amount of product you collect
209
What is the percentage yield equation
Actual yield
——————. X 100
Theoretical yield
210
What is atom economy
Measure of the proportion of reactant to atoms that become part of the desired product in the balenced chemical equation
211
Atom economy equation
Molecular mass of desired products
————————————————— X 100
Sum of molecular masses of all products
212
When is the atom economy always 100%
In an addition reaction
The reactants combine to form a single product
213
Why is low atom economy bad
Lots of waste
Costs money to separate desired prices form the waste product and more money to dispose of the water products safely so they don’t harm the environment
Reactant chemicals are usually costly so it’s a waste of money if a high proportion end up as useless products
214
Why are reaction with low atom economy less sustainable
Many raw materials are in limited supply, so it makes sense to use them efficiently so they last as long as possible.
215
Apart form atom economy and % yield what else is considered in making a reaction sustainable
Temp and pressure
Lower temps and pressure are cheaper and better for the environment
Raw materials that come from renewable sources are better for the environment,ny that materials from non renewable sources
216
Why might a reaction not give a 100% yield
1) some reactions may not go to completion (not finished / reversible)
2) some products may be lost when it is separated for the reaction mixture
3) unexpected reaction might occur
217
What is a standard solution
A solution of a known concentration made by dissolving a primary standard in a suitable solvent
218
What is a concentrated solution
A solution that has a high concentration of solute
219
What is a primary standard
a soluble solid compound that is very pure, with a consistent formula that does not change on exposure to the atmosphere, and has a relatively high molar mass.
220
In order for a solid to be suitable for a primary standard a compound must…?
· be readily available (and inexpensive)
· be very pure (analytical reagent grade, 99.9% pure)
· have a known formula (e.g. degree of hydration must not change)
· be unaltered in air during weighing, that is, the compound must
NOT(a) absorb moisture from the air
(b) absorb carbon dioxide from the air
(c) be oxidised by air
· have a relatively large molar mass (to minimise errors during
weighing)
· be soluble (in the solvent used, and under the conditions used, to
make the solution)
221
Why are most alkali hydroxides not suitable for use as a primary standard
· readily absorb moisture, H2O, from the atmosphere
· readily absorb carbon dioxide, CO2, from the atmosphere
222
Whya re HCL and sulfuric acid not suitable for use as a primary standard
although they are both commercially available as concentrated solutions that are easily diluted, the concentration of the "concentrated" solution is NOT accurately known.
223
Why is nitric acid not suitable for a primary standard
Always contains a little nitrous acid, which has destructive action on many acid-base indicators
224
What are some standards that are suitable as primary standards for acid-base tritiations
· anhydrous sodium hydrogen carbonate, NaHCO3, a weak base (we will use this one)
· potassium hydrogen phthalate, KH(C8H4O4), a weak acid
225
What is the method to make a standard solution
1. Weigh by difference 2.65g of Na2CO3 into a beaker
2. Add a small amount of distilled water to the beaker containing the Na2CO3 to dissolve it, stirring with a glass rod.
3. Rinse weighting boat with distilled water and add the washing to the boat
4. Use a funnel boat with distilled water and add the washing to the Beaker
5. Rinse beaker + glass rod with distilled water + add the washing to the flask
6. Make flask up to the graduation, made with distilled water (bottom of miscues on the line)
7. Add stopped and shake the flask by inverting 15 times
226
What are acids
Proton donors.
When mixed with water, all acids release hydrogen ions - H+ (these are just protons, but you never get them by themselves in water they always combine with h20 to form, h30+)
227
What are alkalis
Bases that are soluble in water
They release OH- ions in solution
228
What do alkalis produce in an aqueous solution
OH-
229
How would u describe the reaction between acid and water and bases and water
Reversible
230
Explain a strong acid
Nearly all the acid will dissociate / ionise in water, and nearly all the H+ ions released
231
Are strong acids reversible,
Yes, the equilibrium lies far on the right
232
Explain strong bases
The forward reaction is favoured so nearly all the bases dissociate in water and lots of OH- ions are released
233
Explain weak acids
Backwards reaction favoured
So only small amounts of the acid will dissociate in water and only a few H+ ions are released
234
Explain weak bases
Ionise only slightly in water
Backwards reaction is favoured so only a s,all amount of base dissociated and only a few OH- Ions are released
235
Give an example of a strong acid
HCL
236
Give an example of a strong base
NaOH
237
Give an example of a weak acid
CH3COOH
Ethanoic acid
238
Give an example of a weak base
NH3
239
Acid + alkalis —->
salt + water
240
Explain how acid + alkali = salt + water
1) it’s the hydrogen ions released by the acid and the hydroxide ions released by the alkali that combine to form water
2) you get a salt when the hydrogen ions in the acid are replaced by metal ions or ammonium ions form the alkali
241
What are salts called from sulfuric acid
Sulfates
242
What are salts called from hcl
Chlorides
243
What are salts called from nitric acid
Nitrates
244
Why is ammonia an exception for neutralisation equation
It doesn’t directly produce ions, but aqueous ammonia is still an alkali. This is because the reaction between ammonia and water produces hydroxide ions, ammonia accepts a hydrogen ion form water molecules forming an ammonium ion and a hydroxide ion. So Ammonia can neutralise acids
245
Metal + acid =
Metal salt + hydrogen
246
Metal oxide + acid =
Salt + water
247
Metal hydroxide + acid =
Slat + water
248
Metal carbonate + acid =
Metal salt + carbon dioxide + water
249
What does the neutralise ionic equation show + what is it
( metal hydroxide + acid)
OH- + H20 -> H20
Shows that a proton is transferred from the acid to the hydroxide ion .
250
Ammonia + acid =
Ammonium salt
251
What do all acids contain
Hydrogen
252
When acids dissolve in water and release 1 H+ ion what is it called
Monoprotic
253
When acids dissolve in water and release 2 H+ ion what is it called
Diprotic
254
Examples of bases
Metal oxides
Hydroxides
Carbonates
Ammonia
255
Base definition
a compound that neutralises an acid to form a salt / a proton acceptor
256
Salt definition
The products of a reaction in which the H+ ions form the acid are replaced by metal or ammonium ion
257
What is electronegativity
An atoms ability to attract the electron pair in a covalent bond
258
What is the most electronegative element
Fluorine
259
What is electronegativity measured on
Pauling scale
260
Describe the trends ofelectronegitivity on the Pauling scale
Increases across periods and decreases down groups ignoring noble gases)
261
What makes a covalent bond polar
When the two atoms have different electro negativities, the bonding electrons are pulled towards the more electronegative atom
262
In apolarbond thedifference inelectronegativity between the two atoms causes what?
A permanent dipole
263
What is a dipole
A difference in charge between the two atoms caused by a shift in electron density in the bond, the greater the difference in electronegativity I the more polar the bond
264
Are covalent bonds in diatomic gases polar? And why?
Nonpolar because the atoms have equal electronegativities and so the electrons are equally attracted to both nuclei
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Are bonds polar when the two elements have similar electronegativities
No
266
What do polarbonds have
Permanent dipoles
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What determines whether or not the molecule will have an overall dipole
The arrangement of polar bonds in a molecule
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If the polar bonds are arranged symmetrically _ what does that mean for the polarity
Dipoles cancel each other out so the molecule has No overall dipole and is non-polar
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What if polar bonds are arranged so they don't cancel each other out (dipole)
The charge is arranged unevenly across the whole molecule, and it will have an overall dipole these are polar
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What are the only bonds that ave purely covalent
Bunds between atoms of a single element/ like diatomic gases
271
Most compounds are somewhere in between the two extremes of ionic covalent what does this mean for the properties
Often got coin + covalent properties
272
What can you use electro negativity to predict
What type y bonding would occur between two atoms the higher the difference in electro negativitys, the more ionic n character the bonding becomes
273
What are intermolecular forces
Forces between molecules
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Are intermolecular forces strong or weak
Weak
Much weaker than ionic, covalent or metallic
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What are the 3 types of intermolecular forces
1) induced dipole-dipole or London (dispersion) forces
2) permanent dipole-dipole interactions
3) hydrogen bonding
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What is the strongest type of intermolecular forces
Hydrogen bonding
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What types of intermolecular forces does van den walls forces refer to
Induced dipole-dipole or London (dispersion) forces
Permanent dipole-dipole interactions
278
What do include dipole-dipole forces cause
All atoms and molecules to be attracted to each other
279
Why are induced dipole-dipole forces found between all atoms and molecules
1) elections in charge clouds are always moving reculy quickly, at any particular moment the electrons in an atom are likely to be more to one side than the other_ at this moment the atom would have A temporary dipole
2) this dipole can cause another temporary induced) dipole in opposite directions on a neighbouring atom the two dipoles are then attracted to each other.
3) the second dipole can cause yet another dipole in a third atom
4) because the electrons are constantly moving - the dipoles are being created and destroyed all the time. Even though the dipoles keep changing, the overall effect is for the atoms to be attracted to each other
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Are all dipole-dipole forces the same strange, and why?
No, larger molecules have larger electron clouds, meaning stronger include dipole-dipole forces, molecules with greater surface areas also have stronger induced dipole-dipole forces because they have A bigger exposed electron cloud.
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Why do liquids with stronger induced dipole-dipole forces will have higher boiling points
When you boil a liquid you need to overcome the intermolecular forces, so that the particles can escape from the liquid surface
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What physical properties does dipole-dipole forces affect
Boiling/ meaning points
Viscosity
283
What holds iodine molecules together in a lattice
Induced dipole-dipole forces.
Iodine atoms are held together in pairs by strong covalent bonds
Then molecules are held together in a molecular lattice arrangement by weak induced dipole-dipole attractions
284
What are permanent dipole-dipole interactions
The positive and negative changes on polar molecules cause weak electrostatic forces of attraction between molecules
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Do permanent dipole-dipole interactions happen in addition to or instead of induced dipole-dipole interactions
In addition to
286
When can hydrogen bonding happen
Can only happen when hydrogen is covalently bonded to fluorine - nitrogen or oxygen
287
Why can hydrogen bonding only happen when its bonded toflorine/ nitrogen or O2
Hydrogen has a high charge density because its so small and fluorine,hydrogen , oxygen are very electronegative the bond is so polarised that a weak bond forms between the hydrogen of ome molecule and a lone pair of electrons on the fluorine, nitrogen or oxygen in another molecule
288
Molecules which have hydrogen bonding usually contain what groups
-OH or -NH
289
How do hydrogen bonds effect the properties of substances
They are soluble In water + have higher boiling and freezing points
290
Why does water, ammonia and hydrogen fluoride generally have higher boiling points if u can compare then with other hydrides in their group
There is extra energy needed to break the hydrogen bonds
291
In ice how are h2o molecules held together
In a lattice by hydrogen bonds
292
Does ice or water have more hydrogen bonds
Ice, as when ice melts hydrogen bonds are broken
293
Is ice or water more dense
Ice is less dense as ice has more hydrogen bonds (which are relatively long)
294
In general what is the main factor that determines the boiling points of a substance
The strength of the induced dipole-dipole forces (unless the molecule can form hydrogen bonds)
295
Explain why the boiling points of the group 7 increase from HCI to HI
Although the permanent dipole-dipole interactions are decreasing, the nu. Of elections in the molecule increases so the strength of the induced dipole-dipole interactions also increases.
296
If you have two molecules with a similar number of electrons then the strength of their induced dipole-dipole interactions will be -- - ----?
Similar
297
If one of the substances las molecules that are more polar than the other. What will it meanfor the permanent dipole -dipole interactions
Stronger permanent dipole-dipole interactions and so a higher boiling point
298
Why do simple covalent compounds have low melting and boiling points
The intermolecular forces that hold together the molecules in simple covalent compounds are weak so don't need much energy to break. So the melting and boiling points are normally low - they are often liquids or gases at room temperature
299
AS intermolecular forces get stronger, melting and boiling points -----
Increase
300
As water is a polar molecule, what does it only tend to dissolve
Other polar substances
301
Why are compounds with hydrogen bonds soluble
They can form hydrogen bonds with water molecules
302
Molecules that only have induced dipole- dipole forcesare soluble or insoluble
Insoluble
303
Why don't simple covalent compounds conduct electricity
They are overall uncharged so they can't conduct electricity
304
To melt or boil a simple covalent compounds _ you have to overcome?
The intermolecular forces that hold the molecule together
You don't need to break the much stronger covalent bonds that hold the atoms together in the molecules
305
What 3 factors effectelectronegativity
The distance between the nucleus and the outermost electrons
Ie. Atomic radius
Sheilding
The nuclear charge
306
Describe how the distance between the nucleus and the outmost electrons ( atomic radius) affects electro negativity
The smaller the atom the closer the nucleus is to the shared outer electrons, so the greater the electronegativity
307
Explain how The distance between the nucleus and the outermost electrons
Ie. Atomic radius is shown as a trend on the periodic table
The electro negativity increases as you go up the group as the closer the electrons are to the nucleus (smaller atomic ractius)
308
Describe now shielding effects electro negativity
The larger the atom the more energy levels between the nucleus and the outer energy level so the lessor the electronegativity
309
Explain how shielding is shown as a trend on the periodic table as effecting Electronegativity
The electronegativity increases as you go up the group as the less shielding
310
Describe how nuclear charge effects nuclear charge
The large the nuclear charge (for a given shielding effect) the greater the electro negativity
311
Explain how the nuclear charge is shown as a trend on the periodic table as effecting Electronegativity
Increasing electronegativity as you go right along the period table as there are more protons in the nucleus so more nuclear charge( number of charges on the nucleus increases) This attracts the bonding pair of electrons more strongly
312
W hat is a polar bond
Covalent bond when there is a separation of charge between one end and the other
313
How do polar bonds arise
When one is more slightly electronegative than the other.
The more electronegative atom will attract the electron pair more than the other. So one end becomes slightly negative and the other becomes slightly positive.
314
What is polarity
The unequal sharing of electrons in a bond- if the electron sharing is equal the bond is described as non -polar
315
Larger molecules may contain polar bonds but why are they not necessarily polar
It depends on the shape
A symmetrical molecule ww not be polar even if the individual bonds are polar. The individual dipoles on the bonds cancel out' due to the symmetrical shape of the molecule.
There is No met dipole its non polar
316
What are intermolecular forces
Wear attractive forces between molecules
317
What is the bond enthalpy for London forces
1-10
318
What is the bond ethalpy for permanent dipole-permanent dipole interactions
3-25
319
What is the bond enthalpy for single covalent bond
IS0-500
320
What is the bond ethalpy for hydrogen bonding
10 - 40
321
What intermolecular forces doall molecules have
London forces
322
What increases the strength of the London forces
More electrons
323
What leads to a permanent dipole
Great difference in electronegativity
324
Where is hydrogen bonding found
Large electro negative atoms
Lone pairs
So mainly oxygen, nitrogen, fluorine
325
What 3 elements do hydrogen bunds form between
Oxygen
Nitrogen
Fluorine
326
Why is water a liquid at room temp.
Hydrogen bonding requires lots of energy to overcome
This increase the melting point and boiling point of water
327
How do London forcesform between molecules
Instantaneous dipoles form due to constant movement of electrons - which causes an induced dipole in a neighbouring molecules. These form London forces between the molecules (weak attractiveforces form)
328
Why are pd-pd forces stronger than London forces
Greater electronegativity difference between bonded atoms form stronger attractive forces
329
Why is ice less dense than water
Ice has an open lattlee, h bonds hold water molecules apart
When ice melts, H20 molecules move closer together
So ice is less dense then water
330
Why does water exist as a liquid at room temperature?
Water has hydrogen bonds in combination with DP-DP and London forces takes more energy to overcome and therfore is a liquid at room temp.
331
What are tritations
Techniques used to accurately measure the volume of one solution that reacts exactly with another solution.
Used for finding the concentration of a solution,. Identifying unknown chemicals, finding the purity of a substance
332
What is needed for valid results in tritations
Precise t accurate measurements needed.
333
Is tritation quantitative or qualitative
Quantitative
334
Describe the method for a tritation
1. Use a pipette and pipette filler to add 25 cm3 of the unknown solution to a clean conical flask.
2. Add a few drops of a suitable indicator and put the conical flask on a white tile.
3. Fill the burette with your standard solution. Flush the tap through to remove any air bubbles. Ensure the burette is vertical. Record the initial volume in the burette to the nearest 0.05cm3.
4. Slowly run the solution from the burette into the conical flask, swirling to mix. (The mixture may at first change colour, and then back again when swirled.)
5. Stop adding the acid when the end-point is reached (when the colour first permanently changes). Note the final volume reading to the nearest 0.05cm3. This is your rough/trial titration to determine your approximate titre.
6. Repeat steps 1 to 5 accurately until two results are repeatable (in close agreement). Ideally these should lie within 0.10 cm3 (concordant) of each other.
335
In a tritation where is the known concentration placed
In the burette
336
In a tritation where do you put the solution of an unknown concentration
In the conical flask
337
What is the volume of solution added from the curette called in a tritation
Titre
338
Name a suitable indicator for a tritation
Phenolphthalein
Methyl orange
339
Describe the colour change in phenolphthalein
Clear in an acid
Pink un an alkali
340
Describe the colour change for methyl orange
Red un acid
Yellow in alkali
341
Why can universal indicator not be used un imitations
Needs to change colours at a specific ph - so can be used to precisely identify when the neutralisation reaction is complete.
342
What are concordant results
Within 0.10 of each other
343
What does a pipette measure
Fixed volumes
344
What does A b urrette measure
Different volumes + lets you aced solutions drop by drop
345
When taking a reacting from a bureette, where should you take the reading from
The bottom of the meniscus
346
To make tritation readings more accurate, what should the readings be to the nearest
0.05 cm3
347
What is the analyte in tritations
Concentration of the unknown solution
348
What is concentration measured in
Mol dm -3
349
How many Cm 3 in a litre
1000
350
What errors could occur during tritation
The burette should be rinsed out with the substance that will be put in it
-If it is not rinsed out, the acid or alkaline added, may be diluted by residual water in the barrettes, or may react with substances left from a previous tritation
-This would lead to the concentration of substances being lower and a larger titre being delivered
-If the Jet space in the burette is not filled properly prior to commencing the tritation it will lead to errors if it then fill during tritation -leading to a larger than expected, titre reading
351
What is concentration
The amount of solute in moles, dissolved in 1dm3
352
What is deliquescent
Absorbed moisture from air
353
What is efflorescent
Loses water of crystallisation
354
What is the equation to find % uncertainty
Percentage uncertainty =
How many times equipment was used x uncertainty
————————————————————————- X 100
Meausred value
355
What do all acids contain
Hydrogen
356
What is a monoprotic acid
If 1+ ion released
357
What is a diprotic acid
If 2+ ions released
358
Give examples of some bases
Metal oxides
Hydroxides
Carbonates
Ammonia
359
Defition of an acid
A species that releases H+ ions in aqueous solution / a proton donor
360
Define an alkali
A type of base that dissolves in water to release OH- ions
361
Define a base
A compound that neutralises an acid to from a salt / a proton acceptor
362
Define a salt
The products of a reaction in which the H+ ions from the acid are replaced by metal or ammonium ions
363
What is oxidising agent
substance which oxidises something else
364
What is a reducing agent
Reduces something else
365
What is oxidation number
Based on a set of rules that apply to atoms and can be thought of as the number of electrons involved in bonding to a differnt elements
366
What is the oxidation number for elements
0
367
How is oxidation number written
Sign placed before the number
Eg,
+2
368
What is the oxidation. Number the same as
However…
The charge
However the sign comes be for the number
369
What is hydrogens oxidation number normally
+1
370
What is oxygens oxidation number usually
-2
371
What do oxidation numbers add to (not on a charged ion )
0
372
How can oxidation be described as in terms of hydrogen
Oxidation is loss of hydrogen
373
How can reduction be described as in terms of hydrogen y
Reduction is gain of hydrogen
374
Oxidising agents themselves are…
Reduced
375
Reducing agents themselves are …
Oxidised
376
What is the oxidation number for nitrite (lll)
+3
377
What does the Roman numeral after a compound / element mean for oxidation number
That’s its oxidation number
Eg, iron (ll) is +2
378
Oxidation number of oxygen us -2 except what
Peroxides
379
The oxidation number of hydrogen is +1 except what
Hydrides
380
Are most elements oxidation states the same
Yes
381
What are group ones oxidation number
+1
382
What is group 2s oxidation number
+2
383
What are groups 7s oxidation number mainly
-1
384
What element in group 7s oxidation number is always -1
Fluorine
385
The oxidation of a simple, monatomic ions is….
:the same as its charge
386
What is rhe oxidation charges in peroxides
-1
387
If an element can have multiple oxidation numbers or isn’t in its normal oxidation state what does it have
Shows using Roman numerals
388
Ions with names ending with -ate
Contain ….
Oxygen and another element
389
Sometimes the other element except oxygen in an -ate ion can exist in differnt oxidation numbers and so forms…
Differnt -ate ions
390
What is the reaction stoichiometry
How many moles of one reactant react with how many moles of another reactant
391
What are polyprotic acids
Donate more than one proton
392
An element is simultaneously oxidised and reduces Is called……
Disproportionation
393
What are the rules to help complex half equations
1) balence atoms
Is you need more oxygen add h20
If you need for hydrogen add H+ ions
2) balence charge
- add electrons
394
What is a redox reaction
Both oxidation and reduction take place to differnt elements
395
An increase in oxidation number =
Oxidation
396
A decrease in oxidation number =
Reduction
397
When metal forms compounds is their oxidation number usually positive or negative
Positive
398
What non metals form compounds do they generally have positive or negative oxidation number s
Negative
399
Metals are,……. when they react with acids
Oxidesed
400
Explain the redox with metal and acid
Metal atoms are oxidised, loosing electrons to from positive metal ions
The hydrogen ions in solution are red used gaining electrons and forming hydrogen molecules
