EL Flashcards
Elements of Life (267 cards)
1
Q
Define an Isotope
A
Atoms of the same element with different numbers of neutrons but the same number of protons
2
Q
What is the equation for relative atomic mass?
A
(Relative abundance x isotopic mass)+(relative abundance x isotopic mass) etc.
———————————————————
100
3
Q
What is avagadros constant?
A
6.02 x1023
4
Q
What is the equation linking moles, mass and Mr?
A
Moles(mol)= mass (g) / Mr
5
Q
Define nuclear fusion
A
Lighter nuclei are fused together to form heavier nuclei, which releases and enormous amount of energy
6
Q
In stars where are the heaviest elements found?
A
Centre
7
Q
Why does a heavyweight star become unstable?
A
When the centre is iron, iron absorbs energy during fusion instead of releasing it
8
Q
What happens to a heavyweight star?
A
Becomes unstable and explodes into a supernova. Disperses elements as gas and dust, restarting the cycle
9
Q
What happens to a lightweight star?
A
Once they run out of hydrogen, the expand into a red giant. Eventually resulting in outer gases drifting away to leave a white dwarf
10
Q
What are the differences between a lightweight and heavyweight star?
A
- Lightweight stars can only do nuclear fusion of hydrogen, they are not as hot, they will last longer
- Heavyweight stars can do nuclear fusion of elements up to iron, they are at higher temperatures and pressures but won’t last as long
11
Q
What is mass spectrometry?
A
A measure of the atomic or molecular mass of different particles to find relative isotope abundances
12
Q
How does mass spectrometry work?
A
Atoms or molecules with be ionised to positively charged cation. This ions are separated according to their mass (m) to charge(z) ratio, m/z
13
Q
What is a mass spectrum?
A
A graph where the relative abundance of each ion can be calculated from the height of each peak
14
Q
How are stars formed?
A
By the culmination of dust and gas originating from the big band
15
Q
What is spectroscopy?
A
The study of how light and matter interact
16
Q
How can we recognise elements in space?
A
Under certain conditions, a substance can absorb or emit electromagnetic radiation in a recognisable way. By analysing this radiation we can recognise and find out information about the substance
17
Q
In the electromagnetic spectrum where is frequency to highest?
A
In the gamma ray area
18
Q
State the electromagnetic spectrum starting and the largest wavelength
A
Radio frequency, microwaves, infrared, visible light, ultraviolet, x-rays, gamma rays
19
Q
How does absorption spectra work?
A
Ions, atoms and molecules will absorb certain frequencies of the emitted radiation. These missing frequencies will show up as black lines and an absorption spectra
20
Q
How is an absorption spectra carried out?
A
White light is passes through a cooler flame, then analysed
21
Q
How does an emission spectra work?
A
When particle absorb radiation they will be raised from their ground state to an excited state. This energy then gets emitted and will appear as a coloured line on a black background
22
Q
Define ground state
A
The lowest energy state
23
Q
Define excited state
A
A higher state of energy
24
Q
Lines ______ at higher frequencies
A
Converge
25
Define continuous spectrum
A seamless transition of colours or wavelengths, white light has a continuous spectrum
26
Light in absorption and emission spectra have a _________
Line spectrum
27
What term refers the the absorption and emission spectra of elements?
Atomic spectra
28
What is the balmer series?
Hydrogen emission spectrum in UV light
29
What is the lyman series?
Hydrogen emission spectrum in visible light
30
Explain wave theory
Light is a form of electromagnetic radiation and behaves like a wave which a characteristic wavelength and frequency
31
What is the speed of light travelling in a vaccuum?
3.00 x10*8* ms*-1*
32
What is the equation connecting speed of light, wavelength and frequency?
Speed of light,c,ms*-1*= wavelength, lambda, m x frequency,mew(v), s*-1*
33
Explain particle theory
Light can be explained as a stream of **photons** the energy of the photons depends on the light’s position in the EM spectrum
- proposed by Einstein in 1905
34
What are **photons**?
Tiny packet of energy
35
State the equation to find the energy of a photon
Energy of photon,E,J= Planck constant,h,JHz x Frequency,mew, s*-1*
36
What is the value of Planck constant?
6.63 x10*-34* JHz or Js*-1*
37
Explain Bohrs theory
Bohr’s theory explains why the hydrogen atom only emits a specific number of certain frequencies. The light that a substance absorbs or emits is specific to itself. When an atom is excited atoms will become excited and when the fall back to ground state, they emit the extra energy or photo as EM radiation. Bohr’s theory explained both emission and absorption spectra
38
The energy of a photon is…..
The difference between two energy levels
39
Why was Bohr’s theory considered controversial?
It relied on the pry of the quantisation of energy
40
An electron can only….
Possess definite quantities of energy
41
Define quanta
Discreta packets of energy
42
If an electron is far away from the nucleus, does it have a higher or lower energy level?
Higher
43
A large gap between energy levels is an energy level diagram mean what?
The frequency is high
44
Why do lines on a spectra converge?
Lines on a spectra converge at higher frequencies as the higher energy levels are closer together
45
What are the similarities of an absorption and emission spectra?
- Each element with have a characteristic atomic spectra
- Both use light to identify elements/atoms
- Lines at the same place and same thickness
- lines converge at higher frequencies
46
What does an emission spectra look like?
Coloured lines on a black background
47
What does an absorption spectra look like?
Black lines on a coloured background
48
How do you choose the appropriate amount of significant figures?
Recognise the least amount of s.f. used in the equation
49
Why do flame tests work?
As energy raises the electrons to an excited state they will drop back down again. The colours are seen as the electrons fall back down to ground state. The colour depends on the wavelength which will be different for every element as the energy difference between energy levels are unique
50
Why are atomic spectra helpful?
- Can be used to identify elements
- Intensity of the lines show abundance
- Provides information about the element’s structure
51
What is each shell labelled by?
A principle quantum number, with the higher numbers referring to the outer shells
52
What is the maximum number of electrons in the 2nd shell?
8
53
What is the maximum number of electrons in the 1st shell?
2
54
What is the maximum number of electrons in the 3rd shell?
18
55
What is the maximum number of electrons in the 4th shell?
32
56
What is the **Aufbau** principle?
Electrons fill the lowest energy shells first
57
What subshells are there, and how many electrons can they hold?
- s-sub-shell - 2 electrons
- p-sub-shell - 6 electrons
- d-sub-shell - 10 electrons
- f-sub-shell - 14 electrons
58
Is the energy of sub-shells fixed?
No, there are different energies per type of sub-shell and the energy of each sub-shell depends on the rise and fall off the charge on the nucleus
59
How many orbitals does a s-sub-shell have?
1 s-orbital
60
How many orbitals does a p-sub-shell have?
3 p-orbitals
61
How many orbitals does a d-sub-shell have?
5 d-orbitals
62
How many orbitals does a f-sub-shell have?
7 f-orbitals
63
What is **Heisenburg's Uncertainty Principle**?
The idea that the position of an electron can't be mapped out exactly so an atomic orbital refers the the 95% probability of finding an electron that region
64
What are atomic orbitals?
A reference to a particular region of space around the nucleus
65
What is the max. number of electrons able to fit in an orbital?
2
66
Electrons can only occupy the same orbital is they have ___________? This is know as the ______________?
Opposite spins (clockwise and anti-clockwise)
Pauli Exclusion Principle
67
What shape in a s-orbital?
Spherical
68
What shape is a p-orbital?
Dumbbell-shaped
69
What is electronic configuration?
The arrangement of electrons in shells and orbitals
70
What is **Hund's principle of maximum multiplicity**? And explain
Orbital are first all filled singly and only fill with two electrons once every orbital has been singly filled. This happens because this arrangement is one which keeps the electrons as far apart as possible.
71
How are orbitals filled?
To produce the lowest energy arrangement possible
72
Electrons in singly occupied shells have __________.
Parallel spins
72
Write the electronic configuration of oxygen
1s*1* 2s*2* 2p*2*
This can be shown with boxes for each orbital and arrows within the boxes to signify electrons. Usually drawn going up the page
73
What elements do not follow the electronic configuration pattern? And why?
Elements in Period 3 up to nickel (besides chromium and copper). This is because The 4s sub-shell has a lower energy level that the 3d sub-shell when empty, but higher when filled. So 4s will fill first and empty first
73
How can electronic configurations be abbreviated?
Put the noble gas symbol and the the extra sub-shells e.g. [Ar] 4s*2* 3d*1*
73
How did Mendeleev arrange the elements?
- In order of increasing atomic mass
- Elements with the same properties were in the same vertical group
- Mendeleev left gaps for elements that he predicted and made predictions about these elements properties
74
Who are the main chemists who contributed to the development of the periodic table?
- Johann Dobereiner
- Lothar Meyer
- John Newlands
- Dmitri Mendeleev
75
How are some elements made synthetically?
Made by bombarding uranium atoms with nuetrons
75
Why was Mendeleev's grouping inaccurate?
Due to the existence of unknown isotopes
76
How is the modern periodic table arranged?
- In order of increasing atomic number
- Based of Mendeleev's model from 1869
- Split into four blocks: s,p,d,f
- Vertical columns known as groups that are based on the electron number in the outer shell, and similar physical properties
- Horizontal rows called periods that are based on the number of the shell being filled
76
Define periodicity
the occurrence of periodic patterns
77
What happens when elements are melted/boiled?
the intermolecular forces between the atoms must be broken. The energy needed to break these bonds depends on the strength of the bonds
78
How are the chemical properties of an element decided?
Decided by the (number of) electrons in the outer shell
78
What is the pattern across periods for m.p. and b.p.?
An initial increase but the fall dramatically
78
Define **closed shell** arrangements
A particularly stable arrangement where sub-shells have been fully occupied by electrons
79
How are elements sorted into s/p/d/f block?
Based off of the current sub-shell being filled by electrons
80
How does atomic radii change the reactivity of an atom?
As the number of shells increase, shielding increases and this leads to the first ionisation enthalpy decreasing (reactivity increases)
81
How does atomic radii change across a period?
Across a period the number of protons in the nucleus increases but the number of shells stays the same. The number of electrons in the outer shell also increase, this increases attraction between the outer shell and the nucleus leading the shell to be held closer to the nucleus and therefore decreasing atomic radii
82
What is shielding?
Shielding refers to the core electrons repelling the outer electrons, which lowers the effective charge of the nucleus on the outer electrons
83
To become stable an atom must what?
Accept or donate an electron/s in order to make a full outer shell and therefore form an ion
84
Define organic species
Molecules that contain not only carbon, but also at least one other element
85
Describe the Miller-Urey experiment
The scientist put methane, ammonia, carbon dioxide and water into a flask. They heated the mixture and subjected it to electrical discharge that would have mimicked the earths early atmosphere. They found that amino acids had been produced meaning that proteins could’ve been made in earths atmosphere or dense clouds in space. This was used to explain the origin of life on earth.
86
How many electrons are linked to the stability of noble gases?
8 electrons besides Helium
87
Define covalent bonds
A bond formed by a shared pair of electrons, usually between two none metals
88
In terms of charges, describe covalent bonding.
The negatively charged electrons are attracted to the positive charges of both nuclei. The attraction overcomes the repulsion.
89
Define bonding pairs
The electron pairs that form the bonds
90
Define lone pairs
Electron pairs that aren't involved in the bonds / form bonds
91
How do instantaneous induced dipole bonds form?
Electrons are always moving so at any instant there can be a concentration of electrons on one side. This causes a difference in charge, leading to an attraction/force between the two atoms.
Can also be known as intermolecular bonds
92
Define a dative covalent bond
A shared pair of electrons that have been supplied by only one of the atoms
Can be shown in dot and cross models as an arrow instead of a line
Also known as coordinate bonds
93
What are covalent **intramolecular** bonds?
The bonds within a molecule that are usually very strong
94
What bonds are weaker? Intramolecular or electrons static attractions(intermolecular)
And what does this mean?
Intermolecular bonds are weaker meaning that little energy is needed to overcome these bonds. This results in low melting and boiling points.
95
True or false: simple molecules can conduct electricity?
False, there are no free electrons/charged particles
96
What is the electron repulsion theory?
As similar charges repel, **the electron pairs repel each other as far apart as possible**
97
How do lone pairs affect the shape of a molecule?
Lone pairs repel more strongly, so will reduce the bond angle by **2.5º**
98
In drawings of molecules, what does the line represent?
The bonds that line on the plane of the paper
99
In drawings of molecules what do solid triangles represent?
The bonds that come towards you
100
What do solid triangles represent in molecules drawings?
Bonds that come towards you
101
What do dotted lines/triangles represent in molecules drawings?
Bonds that go away from you
102
What is the octet rule
Atoms should have a complete outer shell of electrons
103
Covalent bonds form to pair up electrons to obey the ________?
Octet rule
104
What is expanding the octet?
Where bonding atoms may have more then 8 electrons in the outer shell
Groups 5-7
105
In group 5 how many covalent bonds can be made and which elements?
3 or 5 bonds, P and As
106
In group 6 how many covalent bonds can be made and which elements?
2,4 and 6 bonds, S, Se and Te
107
In group 7 how many covalent bonds can be made and which elements?
1,3,5 or 7, Cl,Br,I and At
108
What is a better rule that the octet rule?
- Unpaired electrons pair up
- The maximum number of electrons that can pair it is the equivalent to the number of electrons in the outer shell
109
Describe a linear molecule
Regions of charge: 2
Bonding pairs: 2
Lone pairs: 0
Bond angle: 180
110
Describe a planar triangular molecule
Regions of charge: 3
Bonding pairs: 3
Lone pairs: 0
Bond angle: 120
111
Describe a tetrahedral molecule
Regions of charge: 4
Bonding pairs: 4
Lone pairs: 0
Bond angle: 109.5
112
Describe a pyramidal molecule
Regions of charge: 4
Bonding pairs: 3
Lone pairs: 1
Bond angle: 109.5-2.5= 107
113
Describe a non linear molecule
Regions of charge: 4
Bonding pairs: 2
Lone pairs: 2
Bond angle: 109.5-5= 104.5
OR
Regions of charge: 3
Bonding pairs: 2
Lone pairs: 1
Bond angle: 120-2.5= 117.5
114
Describe a bipyramidal molecule
Regions of charge: 5
Bonding pairs: 5
Lone pairs: 0
Bond angle: 120 AND 90
115
Describe an octahedral molecule
Regions of charge: 6
Bonding pairs: 6
Lone pairs: 0
Bond angle: 90
116
What are major constituent elements?
The element that make up most of your body e.g. oxygen, carbon
117
What is a trace element?
An element that is found in traces in the body e.g. calcium, magnesium
118
What is an ultra trace element?
An element that is found in minor quantities in the body
119
Define relative atomic mass
The relative mass of an atom of an element relative to carbon-12. They have no units
120
What is molar mass?
The mass of one mole, will be equivalent to relative atomic mass
121
What is the equation linking moles, mass and molar mass?
Amount of moles,n,mol = mass,m,g / molar mass , M, gmol*-1*
122
What is relative formula mass?
The mass of all atoms in a compound, Mr
123
What is empirical formula?
The chemical formula showing the simplest ration of the formula units / elements
124
What is the molecular formula?
The chemical formula showing the actual number of of atoms in each element/compound
125
What is the formula unit for a diatomic element?
A molecule
126
What is the formula unit of an ionic compound?
A group of ions
127
What is avagadros constant?
The number of formula units in a mole, 6.02 x10*23*
Named for Amedeo Avogadro
128
What is theoretical yield?
The expected amount of products from a reaction
129
What is experimental yield?
The amount of products actually made from the experiment
130
What factors can reduce yield?
- loss of products from reaction vessels
- side reactions occurring
- impurities in the reactants
- changes in temperature and pressure
- the reaction is an equilibrium reaction
131
What is the equation for percentage yield?
Percentage yield = experimental yield / theoretical yield X100
132
Define the mole
A unit that measures amount of substances in such a way that equal amounts of elements contain equal amounts of atoms
133
Define water of crystallisation
Water molecules that are fitted within the ionic lattice/crystalline compound in a regular manner
134
What are cations?
Positively charged ions
135
What are anions?
Negativity charged ions
136
What are complex ions?
An ion that contains more than one type of atom and consists of a covalently bonded group
137
Acid + alkali/base ———> ?
Salt + water
138
Acid + carbonate ———-> ?
Salt + water + carbon dioxide
139
Acid + metal ———> ?
Salt + hydrogen
140
In an acid + carbonate reaction what can be observed?
Effervescence
141
Water molecules _______ so that ______ charges are _____ to the charge of the ion
Arrange, delta, opposite
142
Whatever are spectator ions?
Ions that aren’t involved the reaction
143
Why are state symbols important?
Shows whether the ionic equation involves precipitation
144
Define precipitation
A suspension of particles produced bye a chemical reaction
145
What is the structure of an ionic compound?
- solid
- giant ionic lattice
- each ions is surrounded by oppositely charged ions
- regularly shaped crystals
146
What are the properties of ionic lattices?
- high melting point, strong electrostatic attractions (stronger attractions for larger compounds)
- conduct electricity when molten or dissolved, ions are free to move
147
What are mobile ions?
Ions that are free to move
148
Define ionic bonds
Oppositely charged ions are formed which are bonded together by electrostatic attraction
149
Define monoprotic
Donates one proton e.g.hydrochloric acid
150
Define diprotic
Donates 2 protons e.g. sulfuric acid
151
Define metallic bonding
The electrostatic attraction between positive metal ions and delocalised electrons
152
What is a giant metallic lattice?
A 3-D structure of positive ions and delocalised electrons, bonded together by strong metallic bonds
153
What’s are the properties of metal?
- ductile
- malleable
-good electrical conductivity
- high melting and boiling points
154
Define delocalised valence electrons
Electrons that are free to move
155
Define a simple molecular lattice
A 3-D structure of atoms bonded together by weak intermolecular forces
156
What are the properties of a simple covalent structure?
- low melting/boiling point
- do not conduct electricity
- soluble in non-polar solvents
157
Define van der Waals’ forces
Instantaneous induced dipoles
158
Define a giant covalent lattice
A 3-D structure of atoms bonded together by strong covalent bonds
Insoluble in polar and non-polar solvents
159
Describe the structure of a diamond
- tetrahedral structure
- poor electrical conductivity
- very hard to
160
Describe the structure of graphite
- **GIANT COVALENT STRUCTURE**
- strong hexagons, structure with weak van der Waals’ forces between layers
- good electrical conductivity
- soft
161
How soluble are nitrates?
All nitrates are soluble
162
How soluble are carbonates?
- ammonia carbonate (NH4)2CO3
- Most group one carbonates
- all other carbonates are insoluble
163
How soluble are sulphates?
- most sulphates are soluble
- BaSO4, PbSO4, SrSO4 are insoluble
164
How soluble are chlorides?
- most chlorides are soluble
- AgCl, PbCl2 are insoluble
165
How soluble are hydroxides?
- group one hydroxides, NH4OH are soluble and group 2 hydroxides are more soluble as you go down the group
- most hydroxides are insoluble
166
How soluble are halides?
- most halide salts are soluble
- all silver halides are insoluble
167
All sodium, potassium and ammonium salts are _____?
Soluble
168
What is the symbol equation for nitrates?
NO3 -
169
What is the symbol equation for carbonates?
CO3 2-
170
What is the symbol equation for sulphates?
SO4 2-
171
What is the symbol equation for chlorides?
Cl-
172
What is the symbol equation for hydroxides?
OH-
173
What is the symbol equation for halides?
Cl-, Br-, I-
174
What group are the alkali metals?
group 1
175
What group are the alkali earth metals?
group 2
176
What two trends are seen in the periodic table regarding how metallic the element is?
- elements become more metallic down a group
- elements get less metallic across a period
177
Define metallic
the ease at which an outer electron is lost
178
S-block metals tend to be ________________?
soft and weak with low melting points
179
Are group 1+2 metals reactive?
Yes, very
180
In a group why are there similarities?
the electron configuration is similar
181
In a group why are there differences?
the size of the atom is increasing down the group
182
Is energy needed for in ionisation?
Energy is always needed to overcome attraction between the electron and the nucleus
183
Define first ionisation enthalpy
the energy required to pull an electron out of an atom
(to form a positive ion)
184
What is plasma?
ionised gas
185
Are group 0 ionisation enthalpies high or low?
high, as they are very unreactive
186
What happens to ionisation enthalpies as you go across a period?
they increase as the size of the atom decreases which increases the strength of the attraction between electron and nucleus
187
What is the general equation for the first ionisation process?
X(g) --> X*+*(g) + e*-*
188
What happens to ionisation enthalpies as you go down a group?
they decrease as the size of the atom is increasing as well as the effect of shielding. this decreases the attraction between electron and nucleus making the electron easier to remove
189
What do ionisation enthalpies prove?
They prove the existence of sub-shells
190
How many successive ionisation enthalpies can one atom have?
As many as the number of electrons available to remove
191
How does ionisation energy change between enthalpies?
each successive enthalpy requires more energy than the last, each IE is bigger than the one before
192
What charge do group 2 ions have?
2+
193
M +2H*2*O ---> ?
metal and water
M(OH)*2* + H*2*
metal hydroxide and hydrogen
194
2M + O*2* ---> ?
metal and oxygen
2MO
metal oxide
195
What happens when a carbonate is heated?
it decomposes to form the oxide releases carbon dioxide
196
What happens to thermal stability down a group?
Thermal stability increases so it is harder to decompose
197
Define thermal stability
the ability of a material to resist the action of heat energy
198
The ______ the ion the ______ the charge density
- smaller, higher
- larger, lower
199
What is the effect of a high charge density?
cations with a high charge density will distort or **polarise** the carbonate anion. this makes the compound less thermally stable and therefore easier to decompose
200
Define charge density
a measure of the concentration of charge on the ions
201
Are metal oxide/hydroxides alkaline or acidic in water?
alkaline, although the are not very soluble
202
Are non-metals alkaline or acidic in water?
acidic
203
Describe the trend in reactivity with water down group 2
increases
204
Describe the trend in thermal stability of carbonate down group 2
decomposes at an increasingly high temp.
205
Describe the trend in pH of hydroxide in water down group 2
increasing pH
206
Describe the trend in solubility of hydroxide down group 2
increasing solubility
207
Describe the trend in solubility of carbonate down group 2
decreasing solubility
208
Define thermal decomposition
the breaking up of a chemical substance with heat into at least two chemical substances
209
Why is each IE bigger than the one before?
- for each electron removed, the repulsion from nucleus will be less so shell is drawn closer to the nucleus
- distance from nucleus decreases so attraction is stronger
210
What is a titration?
a **quantative procedure** to find an unknown in a solution by reacting it with a standard solution
211
What is a standard solution?
a solution that has a known (concentration)
212
Atoms are not ________ or________, they are ___________ from one___________ to ________ / simply _________
created, destroyed, only transferred, form, another, rearranged
213
Describe an acid
- proton donor
- turns litmus red
- neutralised by bases
- release H+ when mixed with water
- an acid that donate 1 electron is known as a **monoprotic** acid
- an acid that donates 2 electrons is known as a **diprotic acid**
214
Define an acid
a compound that dissociates in water to produce hydrogen ions, H+
215
Give examples and symbols for some acids
- hydrochloric acid HCL
- nitric acid HNO3
- sulfuric acid H2SO4
- ethanoic acid CH3COOH
216
Describe and define a base
- known as a proton acceptor
- a compound that reacts with an acid to produce water
217
Define an alkali?
a base that dissolves in water to form hydroxide ions, OH-
218
What is the **Bronsted-Lowry theory of acids and bases**?
the theory of H+ transfer
219
Describe oxonium
- H3O+(aq)
- acts as an acid
- often gets shortened to H+(aq)
220
Neutralisation reactions refer to _________
an acid and alkali reacting
221
Write the ionic equation for a neutralisation reaction
H+(aq) + OH-(aq) ---> H2O(l)
222
How are insoluble salts formed?
in precipitation reactions
223
What must you be able to do in order to find the **exact reacting volume at neutralisation**?
- dilute a solution
- make up a standard solution
- carry out an acid-base titration
224
Define concentration
the amount of solute, in mol, dissolved per 1dm*3* of a solution
225
What is the equation for converting concentration in grams to the preferred concentration in moles?
concentration,moldm*-3* = concentration,gdm*-3*
---------------------------------
molar mass, gmol*-1*
226
A concentration will depend on...?
- the amount of solute
- thee final volume of the solution
227
What is 1dm*3* equivalent to?
- 1000cm*3*
- 1 litre
228
Write the equation linking moles, volume and concentration
concentration,c,moldm*-3* = amount,n,mol
------------------------
volume,V,dm*3*
229
How else can moldm*-3* bewritten?
M
230
What constitutes as a concentrated solution?
greater than 10 moldm*-3*
231
What constitutes as a dilute solution?
normal bench solution, 1-2 moldm*-3*
232
Acid + alkali ----> ?
salt + water
- neutralisation
233
Acid + metal ---> ?
salt +hydrogen
- redox
234
Define concordant titres
titres that lie within 0.10cm*3* or less of each other
235
How do you calculate percentage error?
% error = precision of instrument used
--------------------------------------------- x100
measurement made
236
How do you calculate the precision of the instrument?
The precision of an instrument is half the marked precision
237
What equation links moles, volume and concentration?
Moles,n,mol= conc.,c,moldm*-3* x volume,v,dm*3*
238
How do you find concentration in gdm*-3*?
Conc. gdm*-3* = mass of solute,g / volume, dm*3*
239
Acid + carbonate —->?
Salt +water +CO2
240
Reduction is …
Gain (of electrons)
241
Oxidation is….
Loss (of electrons)
242
Describe the properties of group 1 metals
- soft
- shiny when cut but quickly tarnish
- very reactive, reactivity increases down the group
- have relatively low melting points, decrease down the group
- have low melting points
- alkali metal
243
Describe the properties of group 2 metals
- reactive, reactivity increases as you down the group
- relatively high m.p. And b.p., decreases as you down the group
- low density
- form colourless compounds
- solubility in water increases down the group
244
Describe the reactions in group 2
- reacts **vigorously** with oxygen
- reduces strongly, reacts with water to form hydroxides
-redox reactions
- group 2 oxides are bases
- group 2 carbonates decrease in solubility down the group
- thermal stability increases down the group
245
Why does reactivity increase down group 1+ 2?
The increase in distance between the outer electron and the nucleus, and the increased electron shielding as you go down the group , far outweigh the increase in nuclear charge
246
How do you carry out a practice titration?
- use the pipette and filler to add 25cm*3* of alkali to clean conical flask on a white tile
- add a few drops of indicator into the flask
- fills burette with acid and note starting volume
- slowly add the acid to the alkali, swirling to mix
- stop adding acid when the end-point is reached, not final volume reading
- repeat steps until you get concordant titres
247
What happens to the trend in ionisation energies in period 3?
The trend in generally upward but falls at aluminium and sulphur
248
Explain the upwards trend in ionisation enthalpies for period 3
The increasing nuclear charge that causes greater attraction. This drags the outer electrons closer
249
Explain the fall in ionisation energy for aluminium
Aluminiums outer electron is in a 3p orbital. This electron is further from the nucleus and is partially screened by the 3s. These factors offset the extra proton effect
250
Explain the fall of ionisation energy at sulphur
The removed electron is one of the 3p pair. The repulsion between the two electrons in the same orbital makes it easier to remove
251
As you go across period 3 the atomic radii ______?
Decreases, however argon is ignored as it has a vander Waals radius
252
What is electronegativity?
The measure of the tendency of an atom to attract to a bonding pair
- argon is not included
- electronegativity increases across the period
253
Define energy level
A fixed distance front the nucleus of an atom where electrons may be found
254
Define quantised
Restricted to certain discrete magnitudes
255
Explain why atomic radii decreases across a period
The number of protons increases by one. But the number of shells stays the same. Therefore electrons are held more strongly to the nucleus
256
What is **the law of conservation of mass**?
Atoms are not created or destroyed in chemical reactions, just simply rearranged
- Antoine Lavoisier 1774
257
What is **stoichiometry**?
It studies the amounts of substances that are involved in a chemical reaction
- relies on the molar quantities in the balanced equation
258
Why is **stoichiometry** useful?
It tells you
- the quantities of reactants needed to produce a requires quantity of product
- the quantities of products produces from a known quantity os reagents
259
True or false, oxonium is present in every solution of acid in water
True
260
Define Amphoteric
Can act as an acid or a base
- H2O
