Elements Of Life Flashcards
1
Q
Moles =
A
Mass divided by mr
2
Q
Percentage yield =
A
Experimental yield divided by actual yield. x 100
3
Q
Concentration =
A
mass or moles divided by volume
4
Q
Avogadros constant
A
6.02 x 10^23 = 1 mole
5
Q
1 mole =
A
Atoms in 12g of carbon
6
Q
Hydrated
A
When water is present
7
Q
Anhydrous
A
When water is removed
8
Q
Water is slightly polarised
A
Because the hydrogen has a slightly positive charge and oxygen has a negative charge
9
Q
Ions are held together
A
By strong electrostatic forced of attraction
10
Q
Group 2 metals in a carbonate, precipitate intensifies
A
Down the group
11
Q
Group 2 metals in a hydroxide, precipitate intensifies
A
Up the group
12
Q
First ionisation enthalpy
A
This is the energy required for an electron to be pulled out of an atom in a gaseous state
13
Q
Sub shells existence is supported by
A
Ionisation enthalpy
14
Q
Group 2 has a greater charge density
A
Distorting the carbonate ion more making it easier to separate
15
Q
Cm3 to dm3
A
Divide by 1000
16
Q
Speed of light =
A
Wavelength x frequency
17
Q
Energy =
A
Planck constant x frequency
18
Q
Acid + metal
A
Salt + hydrogen
19
Q
Acid + metal oxide
A
Salt + water
20
Q
Acid + metal hydroxide
A
Salt + water
21
Q
Acid + metal carbonate
A
Salt + water + carbon dioxide
22
Q
Abundance of isotopes =
A
(% x Mr) + (% x Mr) all divided by 100
23
Q
Relative abundance =
A
Relative intensity divided by total relative intensity all x 100
24
Q
Acid
A
Has a ph less than 7
A compound that dissociates in water to produce H ions
A proton donor
25
Alkali
Has a ph greater than 7
| A base that dissolves in water to produce OH ions
26
Base
A compound that reacts with an acid to produce water
| A proton acceptor
27
Charge density
This is a measure of the concentration of charge on an ion
28
Empirical formula
The empirical formula tells you the simplest ratio of the different atoms in a compound
29
Group
A vertical column in the periodic table
30
Molar mass
The molar mass of a substance is the mass of 1 mole of it
31
Mole
A mole of a substance is 6.02 x 10^23 particles of it
32
Molecular formula
Tells you the actual numbers of the different atoms in a compound
33
Neutralisation
When an acid reacts with an alkali to form salt + water
34
Oxonium ion
Present in every acidic solution
| Formed when a proton is donated to a water molecule and forms H3O ion
35
pH
This indicates how strongly acidic or alkaline a solution is
36
Polarise
The ability of an ion to distort the charge cloud of an oppositely charged ion
37
Relative atomic mass
This is the mass of the formula of a compound relative to an atom of carbon - 12
38
Thermal decomposition
This is the breaking down of a compound using heat
39
Thermal stability
A compound with the greatest thermal stability is the one which needs the highest temperature to decompose it
40
Water of crystallisation
Water molecules fitted in a regular pattern within the crystal lattice of an ionic solid
41
Absorption spectrum
This is produced when electrons move from a lower energy level to a higher one
It looks like a rainbow with black lines on it
42
Alpha radiation
Composed of 2 protons and 2 neutrons
43
Anion
An ion with a negative charge
44
Atomic number
Tells you the benumbed of protons I’m the nucleus
45
Atomic orbital
Sub shells are split into atomic orbitals
46
Cation
An ion with a positive charge
47
Chromosphere
The region outside a stars surface which is made of atoms, ions and molecules
48
Closed shell arrangements
These are typical of noble gases
| Their shells and sub-shells are fully occupied by electrons
49
Covalent bonds
These are formed between two non-metal atoms
| The electrostatic attraction between both nuclei and the pair of electrons holds the atoms together
50
Dative covalent bond
A covalent bond in which both the shared pair of electrons come from the same atom
51
Delocalised electrons
These are seen in metallic bonding
| They do not belong to specific atoms but are shared
52
Electron configuration
This is the arrangement of electrons in shells, sub shells and atomic orbitals
53
Electrostatic bond
It is an attraction between something with a positive charge and something with a negative charge
54
Emission spectrum
This is produced when electrons move from a higher to a lower energy level
It is coloured lines on a black background
55
Excited
An electron is excited if it moves from a lower to a higher energy level
56
Frequency
This is the number of vibrations per second
| Units are Hertz (Hz) or s^-1
57
Planck constant
6.6.3 x 10^-34 Js
58
Fusion
The joining of 2 small nuclei to make a larger one
It occurs at a high speed to overcome repulsion between positive nuclei
High temperature and pressure are needed
59
Giant covalent network structure
This is typical of group 4 elements and their compounds
All of the bonds are strong covalent ones
There are no distinct molecules
60
Giant ionic structure
This is typical of ionic compounds
| The ions are held in a 3D lattice by ionic bonds
61
Giant metallic structure
This is typical of metals
| The metal ions are attracted to the delocalised electrons
62
Group when used to describe electrons
A group can be a single, double or triple bond or a lone pair of electrons
63
Intermolecular bond
Bond between atoms in a molecule
64
Intramolecular bond
Bond between molecules
65
Relative atomic mass
The average mass of an atom of an element on a scale where an atom of carbon 12 is 12
66
Relative isotopic mass
The mass of an atom of an isotope of an element on a scale where an atom of carbon 12 is 12
67
Relative formula mass (Mr)
The average mass of a molecule of formula unit on a scale wearing atom of carbon 12 is 12
68
To find the Mr
Add up The relative atomic mass values of all the atoms in the molecule
69
Relative atomic mass is an
Average of the isotopes so it is not usually a whole number (eg Cl 35.5)
70
Proton charge
+1
71
Neutron charge
0
72
Electron charge
-1
73
Proton relative mass
1
74
Neutron relative mass
1
75
Electron relative mass
1/2000
76
Mass number
The total number of protons and neutrons in the nucleus
77
Atomic number
The number of protons in the nucleus which is also the same as the number of electrons in the nucleus
78
Negative ions have more
Electrons than protons
79
Positive ions have fewer
Electrons than protons
80
Isotopes are
Atoms of the same element with different numbers of neutrons
81
Isotopes have the same
Configuration of electrons so they’ve got the same chemical properties
82
Isotopes have different
Physical properties because physical properties depend more on the mass of an atom
83
John Dalton
Set an atom was a solid sphere and that different spheres made up different elements
84
JJ Thompson
Said the atoms were not solid and came up with the plum pudding model
85
Plum pudding model
Had a positively charged sphere with negative electrons in bedded in it
86
Rutherford experiment
Gold foil experiment and fired alpha particles add a thin sheet of gold most passed straight through and a few deflected so the plum pudding model was wrong
87
Ernest Rutherford did
The gold foil experiment and proved the Plum pudding model was wrong and came up with the nuclear model of an atom
88
Nuclear model of an atom
Small positive charge in the nucleus
Nucleus surrounded by a cloud of negative electrons
Mostly made up of empty space
89
Mosley discovered
That the charge of the nucleus increased from one element to another
90
Rutherford discovered
That the nucleus contains positively charged particles that he called protons
the charges of the nuclei of different atoms could be explained the atoms of different elements have a different number of protons in the nucleus
91
James Chadwick discovered
But there were other particles in the nucleus that had mass but no charge (neutron)
92
Bohr model
Electrons can only exist in fixed orbitals or shells not anywhere in between
Each shell has a fixed energy
93
When an electron moves between shells
Electromagnetic radiation is admitted or absorbed because the energy of shells is fixed the radiation will have a fixed frequency
94
Noble gases are stable
Because they have a full shell of electrons
95
Quantum model
You can never know where an electron is or which direction it’s going in at any moment but you can say how likely it is to be at any particular point in the atom (The denser the dots the more likely an electron is to be there)
96
Measured relative masses
Use a mass spectrometer
97
Mass spectrometer steps
Vaporisation
Ionisation
Acceleration
Detection
98
Vaporisation
The sample is turned into a gas using an electrical heater
99
Ionisation
The gas particles are bombarded With high energy electrons to ionised them. Electrons are knocked of the particles leaving positive ions
100
Acceleration
The positive ions are accelerated by an electric field
101
Detection
The time taken for the positive ions to reach the detector is measured this depends on an iron mass and charge (light highly charged ions will reach the detector first while heavy ions with a smaller charge will take longer)
For each sample analysed a mass spectrum is produced
102
A mass spectrum
Y axis gives the abundance of the ions whilst the X axis gives the mass overcharge ratio (relative mass)
103
Mass spectrum recognising
Isotopes- if the sample is an element each line will represent a different isotope of the element
104
To find the AR using a mass spectrum
For each peak read the relative isotopic abundance and times by the relative isotopic mass
At the totals and then divide by 100
105
Percentage relative isotopic abundance =
(Relative abundance/ total relative abundance) x 100
106
To find the Mr using a mass spectrometer
The peak further to the right shows the Mr
107
Fragmentation pattern
Bombarding with electrons make some molecules break into fragments these are shown on a mass spectrum
108
If an atom is unstable
It will break down to become stable, the instability Could be caused by having too many neutrons or not enough neutrons or too much energy in the nucleus
109
Breaking down of an unstable atom is called
Radioactive decay
110
Alpha particles
Like helium I stopped by paper and have a strong ionising ability and slight deflection in electric field
111
Beta particles
A fast moving electrons stopped by thin aluminium sheets they have a moderate ionising ability and a large deflection in electric field
112
Gamma rays
A very short wave Electromagnetic waves they are stopped by very thick lead and have a weak ionising ability and then not deflected in an electric field
113
Alpha particles are strongly
Positive so they can remove electrons from atoms
When an alpha particle hits an atom it transfers some of his energy to the atom
The alpha particle quickly ionises lots of atoms and loses all its energy and that is why it has a low penetrating power
114
Beta particles have lower charges but higher speeds
They can still knock electrons of atoms but they hit atoms less frequently because they’re smaller so they have a better penetrating power
115
Nuclear equations
Balance the top and bottom numbers in the equations
14. 14. 0
C N. = e
6. 7. -1
116
Radioactive decay is
Random but for radioactive atoms the pattern is best described using the idea of half life
117
Half life
The average time taken for half of the atoms in a sample to decay, it has a constant value for any particular isotope
118
Radio active isotopes can be used as
Medical traces as it is easy to detect radiation given out
119
Only isotopes are suitable half lives can be used as
Medical traces because a very long half life is dangerous as they are exposed to too much radiation but a very short half life is inconvenient
Alpha emitters are no good as they cause damage by ionising atoms inside the body and they wouldn’t be detectable outside the body
120
Radiocarbon dating
Measures how much of a particular isotope of carbon there is in a plant or animal this can be used to determine the age of rocks and archaeological finds
121
The last carbon-14 in a sample of organic material
The older it must be
122
Hydrogen nuclei combine to
Helium nuclei releasing large amounts of energy through nuclear fusion
123
When the hydrogen in a star is called runs out the temperature and pressure of the core starts
To rise and it’s all get hot enough to fuse heavier elements
124
Number of moles=
Number of particles you have/ number of particles in a mole (Avogadros constant)
125
Molar mass units
g mol^-1
126
Calculate the mass of….if ...g of ... is burnt in air
Find molar mass
Number of miles
Mole ratio
Moles x molar mass
127
Empirical formula
Gives a smallest whole Number ratio Of atoms in a compound
128
Molecular formula
Guess they have a number of atoms in a molecule
129
Empirical formula is a calculated
Find number of moles (mass/mr)
Mole ratio
Divide both numbers by the smallest mole ratio
C:H 1:2 mole ratio = CH2
130
To work out the molecular formula
At the masses of the empirical formula and see if they equal the molar mass = molecular formula
131
How many electrons can shells hold
2
8
8
16
132
Electrons move around the nucleus in
Shells or energy levels
133
Atoms in the ground state have all their electrons at their
Lowest possible energy levels
134
If an atom is electrons take in energy from the surroundings they can move
To a higher energy level further from the nucleus, the electrons become excited
135
Electrons can release energy by
Dropping to a lower energy level
136
The energy levels all have certain
Fixed values. Electrons can jump from one energy level to another by observing or releasing a fixed amount of energy
137
When electromagnetic radiation is passed through a gaseous element the electrons only absorb
Certain frequency is corresponding to differences between the energy levels
138
Absorption spectrum
Coloured background and back lines
139
Emission spectrum
Black background and coloured lines
140
Absorption spectrum is when
An electron moves to a higher energy level and the electrons becomes excited
141
Emission spectrum is seen when
Electrons dropped down to lower energy levels and they give out certain amounts of energy
142
🔺E =
H x v
143
The difference in energy between two shells (J) =
Frequency (Hz) x Plancks constant (Js)
144
Each line in a spectra represents
Electrons moving to or from a different energy level
145
As the energy levels get closer together
The energy/frequency increases
146
Ionic bonding
Is when ions stick together by electrostatic attraction they are formed when electrons are transferred from one atom to another
147
Compound
When different elements join and bond together
148
Elements in the same group all have
The same number of outer electrons
149
Electrostatic attraction
Holds positive and negative ions together and is often very strong
150
How do you show ionic bonding
Dot And cross diagram
151
Sodium chloride has a
Giant ionic lattice structure
152
A lattice is
Cube shaped
153
Ionic compounds conduct
Electricity when they’re molten or dissolved. As the ions are free to move where is in a solid at the fixed in position by strong ionic bonds
154
Ionic compounds have high
Melting points as they are held together by strong electrostatic forces
155
Ionic compounds are often
Dissolved in water, as water is a polar molecule and pulls the ions away from the lattice
156
Ionic bonding is between
A metal and a non-metal
157
Covalent bonding is between
Two nonmetals
158
Metallic bonding is between
TWO metals
159
Date Covalent bonding is where
Both electrons come from one atom
| So one atom donates both electrons to a bond
160
Molecular substances usually have a fairly low
Melting and boiling point as there is no giant structure that needs to be broken down
161
To Melt or boil a simple molecular compound you only have to overcome
The attractions between the molecules these are pretty weak compared to ionic covalent bonds
162
Molecular substances don’t
Conduct electricity because there are no charge carriers that are free to move
163
Molecular substances are usually
Insoluble in water
164
Ionic bonds are usually in the state
Solid
165
Simple covalent bonds are usually in the state
Liquid or gas
166
Giant covalent structures examples
Diamond which is made of carbon and silicone dioxide which is made of silicone
167
Giant covalent structures have very high
Melting points as you need to break a lot of very strong bonds before they melt
168
Giant covalent structures are often extremely
Hard
169
Giant covalent structures are good
Thermal conductors
170
Giant covalent structures won’t
Dissolve so they are in soluble in polar solvents like water
171
Giant covalent structures can’t
Conduct electricity
172
In metallic lattice is the electrons in the outer most shell of the metal atoms
Are delocalised- this leaves a positive metal ions and a sea of delocalised electrons around it
173
In metallic bonding The positive metal ions are attracted to the
Delocalised negative electrons
174
Metallic bonding has a high
Melting point because of the strong metallic bonds
175
The more delocalised electrons per atom the stronger
The bonding will be in the higher the melting point
176
As there are no bonds holding specific ions together the metal ions can
Slide over each other when the structure is palled and an example of this is graphite which is very ductile
177
The delocalised Electrons can pass kinetic
Energy to each other making metals good thermal conductors
178
Metals are good electrical
Conductors because the delocalised electrons can carry a current
179
Metals are in
Insoluble
180
Electrons are all negatively charged so they will
Repel each other as much as they can
181
Lone pairs of electrons repel
More than bonding pairs
182
The greatest angles are between
Lone pairs of electrons
183
Bonding pair to a bonding pair bond angles
Are the smallest angles
184
Electron pair repulsion principal
Is well electrons repel each other as much as they can with the lone pairs repelling more than the bonding pairs
185
No lone pairs has a bond angle
Of 109.5
186
One lone pair has a bonding angle of
107
187
To lone pairs has a bond angle of
104.5
188
A linear molecule has a bone angle of
180°
189
Three electron pairs on a central atom has a bond angle of
120°
190
A non-linear shape with a lone pair of electrons forms a
Bent shape
191
For electron pairs on a central atom with no lone pairs forms a
Tetrahedral shape
192
For electron pairs on a central atom with a one lone pair forms a
Trigonal pyramid
193
Six electron pairs on a central atom forms a
Octahedral shape
194
In the 1800s there was only two ways to categorise elements
By the physical and chemical properties and by the relative atomic mass
195
John Newlands
Arrange the elements in order of mass, similar elements appeared at regular intervals every eight element was similar this was called the law of octaves
196
Mendeleev
Left some gaps in the periodic table where elements didn’t seem to fit, He was then able to predict the properties of the missing elements by comparison with other elements in the same group
197
The modern periodic table
Arranges elements by proton numbers
198
All the elements within a period
Have the same number of electrons shells
199
All the elements within a group have the same
Number of electrons in the outer shell
200
Properties often change gradually as you go
Down each group
201
For metal is melting and boiling points increase
Across the period because the metal metal bonds become stronger because the metal ions have an increasing number of delocalised electrons and a decreasing radius this leads to a higher charge density which attracts the ions together more strongly
202
Intermolecular forces are weak
And easily broken so these elements have low melting and boiling points
203
More atoms in a molecule mean stronger
Intermolecular forces
204
The noble gases have the lowest
Melting and boiling points because they exist as individual atoms resulting in very weak intermolecular forces
205
Group 2 elements react with
Water to produce hydroxides
206
Group 2 elements get increasingly
Reactive down the group because the outermost electrons are furthest from the nucleus and so more easily lost
207
Group to oxides and hydroxides are
Bases
208
The oxides of the group to metals react readily with water to form
Metal hydroxide which dissolve
209
If the solution contains hydroxide ions the solution is strongly
Alkaline
210
The oxides for more strongly alkaline solutions as you go
Down the group because the hydroxide to get more soluble
211
Group 2 elements that contain single charge negative ions (OH-)
Increase in solubility down the group
212
Group 2 elements that contain doubly charged negative ions (CO3^2- and SO4^2-)
Decrease in solubility down the group
213
Group 2 carbonates decompose
To form the oxide and carbon dioxide
214
The more thermally stable a substance is the more
Heat it will take to break it down
215
Thermal stability increases
Down the group
216
Carbonate ions are a large onions and can be made unstable by the presence of a
Cation
217
A cation draws the electron on the carbonate ion towards itself (polarises it) this then
Distorts the carbonate ion and the greater the distortion the less stable the carbonate ion is
218
Large cations cause less
Distortion
219
The further down the group 2 the larger the
Cations and the less distortion cause and the more stable the carbonate anion is
