Unit 1: Everything Flashcards
1
Q
Elements:
A
-chemical substances that cannot be broken down into simpler substances
-due to the presence of isotopes, not every atom in the element will be identical
-however all atoms in an element will contain the same number of protons (atomic number)
e.g. He
2
Q
Molecules:
A
-molecules are two or more atoms that are bonded together
-molecules can contain atoms of the same element (diatomic element) or two or more atoms of different elements (compound)
e.g. Br-Br
3
Q
Compounds:
A
-compounds are substances containing atoms of different elements combined together
-compounds can also be formed of ions of opposite charges
e.g. H-Br
4
Q
Ions:
A
-ions are substances which have lost or gained electrons to have a charge
-cations are positive ions- atoms that have lost electrons
-anions are negative ions- atoms that have gained electrons
-O-C=O
-O-
5
Q
Other terms:
A
-monoatomic substances are made up of one single atom
e.g. Noble gases
-diatomic substances are made up of two atoms
e.g. N2 or O2
-elements or compounds made up of several atoms are known as polyatomic
6
Q
Reactions of acids: with metals
A
Metal + acid -> salt + hydrogen
e.g. Mg(s) + HCl(aq) -> MgCl2(aq) + H2(g)
7
Q
Reactions of acids: with metal oxides and insoluble metal hydroxides
A
-Metal oxide + acid -> salt + water
e.g. CuO(s) + H2SO4(aq) -> CuSO4(aq) + H2O(l)
-Metal hydroxide + acid -> salt + water
e.g. Zn(OH)2(s) + H2SO4(aq) -> ZnSO4(aq) + H2O(l)
8
Q
Reactions of acids: with alkalis
A
Alkali + acid -> salt + water
e.g. 3NaOH(aq) + H3PO4(aq) -> Na3PO4(aq) + H2O(l)
9
Q
Reactions of acids: with metal carbonates
A
Metal carbonates + acid -> salt + water + carbon dioxide
e.g. Li2CO3(s) + 2HCl(aq) -> 2LiCl(aq) + H2O(l) + CO2(g)
10
Q
Reactions of acids: with hydrogencarbonates
A
Hydrogencarbonate + acid -> salt + water + carbon dioxide
e.g. NaHCO3(s) + HNO3(aq) -> NaNO3(aq) + H2O(l) + CO2(g)
11
Q
Displacement reactions:
A
-displacement reactions occur when one element is replaced with another element in a compound
e.g. redox
-displacement can occur with metals in aqueous solutions, metals in the solid state and halogens
12
Q
Displacement reactions: Metals
A
-Aqueous solution:
Mg(s) + CuSO4(aq) -> MgSO4(aq) + Cu(s)
-this reaction will occur without the need for energy to be added and is typically carried out on a small scale in a lab
-Solid state:
2Al(s) + Fe2O3(s) -> Al2O3(s) + 2Fe(s)
-this reaction requires a very high temperature to start and is done for a specific purpose in industry
-this is called the thermite reaction and is used to weld railway lines together
13
Q
Displacement reactions: copper
A
-when a copper wire is placed in a solution of silver nitrate the silver is displaced
-this can be seen by the build up of silver metal on the copper metal rod
-the colour of the silver nitrate solution also changes from colourless to blue
14
Q
Displacement reactions: thermite
A
-the thermite reaction is another example of a displacement reaction
-aluminium and iron(III) oxide is positiones above the place where the two rails are to be joined
-a magnesium fuse is lit and the molten iron formed fuses the rails together
15
Q
Displacement reactions: Halogens
A
-more reactive halogens can displace less reactive halogens from their compounds
e.g Cl2(aq) + 2KBr(aq) -> 2KCl(aq) + Br2(aq)
-electrons are transferred from the Br- ions to the Cl atoms so bromide ions are oxidised to bromine and chlorine is reduced to chloride ions
16
Q
Precipitation Reactions:
A
-precipitations reactions can be used in chemical tests or to determine the coefficients in a chemical equation
-CO2 gas is detected by adding limewater, Ca(OH)2 and forming a white precipitate of calcium carbonate, CaCO3
-SO42- ions are detected by adding barium chloride, BaCl2, and forming a white precipitate of barium sulfate, BaSO4
-Halide ions are detected by adding silver nitrate, AgNO3, and forming precipitates of silver halides, AgX
17
Q
relative atomic mass (Ar):
A
-> is the weighed average mass of an atom compared to 1/12 the mass of an atom of carbon-12
18
Q
relative molecular mass (Mr):
A
-> is the sum of the relative atomic masses of each atom in a molecule
19
Q
relative formula mass (Mr):
A
-> is the sum of the relative atomic masses of each atom in a formula (including molecules and ions)
20
Q
molar mass (M):
A
-> is the mass of one mole of a substance
21
Q
Avogrado’s constant:
A
-> is the scaling up factor from atoms, molecules and ions to grams
-the value is 6.02 x 10^23mol^-1 - this is the number of particles in one mole of any substance
number of particles = moles (mol) x Avogrado’s number (6.02 x 10^23)
22
Q
mole:
A
-we use a fixed amount called a mole to compare molecules
-one mole of any substance will contain the same number of atoms or molecules = 6.02 x 10^23
moles = mass(g)/molar mass(g mol^-1)
23
Q
Reacting masses calculations: 4 steps
A
1) write a balanced chemical equation
2) determine the number of moles of the known substance
3) use the mole ratio to determine the number of moles of the unkown substance
4) determine the mass of the unknown substance
24
Q
excess calculations:
A
-if a reagant is in excess, we have more than we need to react with the other reactant
-the other reactant is known as the limiting reagent
A + B -> C + D
A = limiting reagent, B = excess
25
percentage yield:
-> is the percent ratio of actual yield to the theoretical yield
-how much product you actually make in an experiment compared to how much you should have made, expressed as a percentage
percentage yield = actual yield/theoretical yield x 100
26
atom economy:
-> is the ratio between the mass of the desired product and the mass of all products formed
-for e.g. a reaction that makes 1 product will have an atom economy of 100% as it only makes one desired product
atom economy = molar mass of desired product/sum of the molar masses of all products x 100
27
formula:
-chemical formulae tell us the number of atoms present in a compound
-for example, SO2 means that there are 2 oxygens of 1 sulfur
-two types:
-empirical
-molecular
28
empirical formula:
-> is the simplest whole number ratio of the atoms present in a compound
-for example: CH2
29
molecular formula:
-> shows that actual number of atoms of each elemnt present in a covalent compound or the formula unit of an ionic compound
-for example: C4H8
30
empirical formula: experimental methods
-the formula of an oxide of copper can be calculated by converting the copper oxide to copper (removing the oxygen):
1. a known mass of copper is placed in a tube and heated in a stream of hydrogen gas
2. the hydrogen reacts with the oxygen in copper oxide to form steam and copper metal
3. the excess gas is burned off, the tube is cooled and weighed
4. the process is repeated until a constant mass is obtained
31
empirical formula: using masses
table with:
-mass of element (g)
-relative atomic mass
-division by Ar
-ratio
32
empirical formula: percentages
table with:
-percentage of element
-relative atomic mass
-division by Ar
-ratio (divide by the smallest moles)
33
empirical formula: combustion analysis
table with:
-mass of an element (g)
-relative atomic mass
-division by Ar
-ratio
34
molecular formula: calculations
-to convert between empirical formula and molecular formula, we must know the Mr of the compound
-determine the Mr of the empirical and divide by Mr of compound to find ratio and then mutiply this by the empirical formula
e.g. C2H5 has an Mr of 29 and the molecular formula is 58, determine the formula -> 58/29 = 2 - 2 x C2H5 = C4H10
35
ideal gas equation:
pV = nRT can be used for gases to find the amount of a substance in mole
-this can then be used to determine the molecular formula from empirical
p = pressure in pascals (Pa)
V = volume in cubic (m^3)
T = temperature in kelvin (K)
n = amount of substance in moles (mol)
R = the gas constant - 8.31 Jmol^-1K^-1
36
converting units:
-kPa -> Pa multiply by 10^3
-cm^3 -> m^3 divide by 10^6 or multiply by 10^-6
-dm^3 -> m^3 divide by 10^3 or multiply by 10^-3
-Cº -> K add 273
37
calculations involving gas volumes:
-Avograd's law states that equal volumes of gases at the same temperature and pressure will have the same number of molecules
-this means that 100 cm^3 of hydrogen will have the same number of molecules of 100 cm^3 of chlorine at the same temp. and pressure
-gases are measured in cm^3, dm^3, or L, 1 L = 1 dm^3 = 1000 cm^3
1 dm^3 -> 1000 cm^3 (multiply by 1000)
38
molar volume calculations:
-> is the volume 1 mole of a gas occupies at room temperature and pressure (rtp)
-is aprox. 24 dm^3 or 24000 cm^3
molar volume (24 dm^3) = volume (dm^3) / moles(mol)
39
concentrations:
-can be measured in g dm^3 or mol dm^-3
mass concentration (g dm^-3) = mass (g) / volume (dm^3)
40
concentrations in PPM:
concentration in ppm = Mass of solute/mass of solvent x 100
41
gas concentrations in PPM:
concentration in ppm = volume of gas/volume or air x 1000000
42
structure of an atom:
-proton, 1, +1 nucleus
-neutron, 1, 0, nucleus
,electron, 1/1840. -1, quantum shells
43
atomic number:
number of protons
44
mass number:
number of protons PLUS the number of neutrons
45
isotopes:
atoms of the same element that have the same atomic number but different mass numbers
-for example:
chlorine-35 has 17 protons and electrons and 18 neutrons
chlorine-37 has 17 protons and electrons and 20 neutrons
46
relative atomic mass:
-> the weighted mean average of an atom of an element compared to 1/12 of the mass of carbon-12
47
relative isotopic mass:
-> the mass of an individual atom of a particular isotope relative to 1/12 of the mass of carbon-12
48
mass spectrometry:
-in a mass spectrometer, particles are turned into positive ions, accelerated and then delfected by an electric or magnetic field
-the resulting path of ions depends on their 'mass to charge' ratio (m/z or m/e)
-particles with a large m/s value are deflected least and those with a low m/s value are deflected most
-the results produe a mass spectrum which portrays the different ions in order of their m/z value
49
uses of mass spectrometry:
-mass spectrometry was initially used to show the identity of isotopes
-it is now used to calculate molecular masses and characterise new compounds
50
how does a mass spectrometer work?
1) vaporisation:
-atoms in a sample are vaporised into a gas
2) ionisation:
-gaseous atoms are bombarded by electrons from and electron gun and are ionised and sufficient energy is given to form ions of +1 charge
3) acceleration:
-ions are charged so can be accelerated by an electric field
4) deflection:
-charged particles will be deflected by a magnetic or electric field
5) detection:
-by electric or photographic methods
51
results of mass spectrometry:
-the radius of the path depends on the value of the mass/charge ratio (m/z)
-ions of heavier isotopes have larger m/z values so follow a large radius curve
-as most ions are +1 charged, the amount of separation depends on their mass
-if an ion acquires a 2+ charge it will be deflected more; its m/z value is halved
52
mass spectrum in diatomic molecules:
-some elements and compounds contain two or more atoms covalently bonded together
-if these substances are analysed by mass spectrometry, you can obtain the relative molecular mass of the element or compound by observing the peaks with the largest m/z ratios (assuming a value of z = 1)
e.g. mass spectrum of chlorine has two peaks:
-m/z 35- 75%
-m/z 37- 25%
it also has 3 peaks: m/z 70, m/z 72, m/z 74
53
quantum shells:
-electrons exist in certain well defined energy levels called quantum shells
-all electrons in a quantum shell have similar, but not identical, energies
-each quantum shell, apart from the first, is further divided into subshells of slightly different energy levels
-the higher the energy of the shell, the more subshells it contains
54
subshells:
-the s orbital is the lowest energy in each shell, followed by the p, then the d subshell and lastly the f subshell
-quantum shell 1 -> s
-quantum shell 2 -> s p
-quantum shell 3 -> s p d
-quantum shell 4 -> s p d f
55
orbitals:
-each subshell has an area occupied by electrons, these areas are known as orbals
-each orbital will hold 2 electrons
-an orbital has no edge: the electron density just tails off more and more
-however, we can calculate the surface which would contain a given percentage (say 95%) of the electron density, and use this to gain an idea of the shape of the orbital
56
s orbitals:
-s orbitals are the lowest energy subshell in each quantum shell
-the 1 s orbital is symmetrically spherical
-the 2s orbital is the same shape just larger
57
p orbitals:
-the 2nd quantum shell conntains three separate p orbitals
-these orbitals have an elongated dumbell shape along one axis- px,py,pz
58
d orbitals:
-the 3rd quantum shell contains five separate d orbitals
-these orbitals have a more complicated shape along two axes- dx2-y2, dz2, dzy, dxz, dyz
59
electrons in subshells:
-each quantum shell is split in to specific subshells and these subshells have an effect on how the electrons are entered into each orbital
-opposte spins!!!
60
electronic configurations:
-electrons are added to the lowest energy orbitals until all electrons are accommodated
-electrons will occupy the orbitals singly before pairing takes place
-two electrons cannot occupy the same orbital unless they have opposite spins
exceptions:
-potassium
-calcium
-chromium
-copper
61
ionisation energy:
-> the energy required to remove one electron from each atom in one mole of gaseous atoms of the element
-this means that electrons in the outer quantum shells (higher energy) will need less energy to be removed in comparison to the electrons in the inner quantum shells
-more than one electron can be removed from an atom so we can also have first, second, third and fourth ionisation energies
61
first ionisation energy:
E (g) -> E+ (g) + e-
metals and non-metals both undergo isonisation to form positively charged ions
62
second ionsation energy:
E+ (g) -> E2+ (g) + e-
63
third ionisation energy:
E2+ (g) -> E3+ (g) + e-
64
shielding:
-electron-electron repulsion is also known as shielding
-this effect raises the energy of the electrons to the above the value they would have if there was no repulsion between them
-the inner electrons cause this effect to the outer electrons
-if an electrons exists in an orbital where it is paired with another electrons it will experience electron-electron repulsion causing a lowe ionisation energy compared to a singly occupied orbital
65
energy of an electron:
-the main factors that affect the energy of an electron are:
-the orbital in which the electron exists
-the nuclear charge of the atom (number of protons)
-the repulsion (shielding) experienced by the electron from the other electrons in the atom
66
trends in ionisation energy:
-as you go across a period there is an increase in the first ionisation energy because of more protons
-as you go down a group the first ionisation energy decreases
67
the periodic table: periods
all elements in a period have the same number of quantum shells containing electrons
68
periodic properties:
-> regularly repeating patterns of atomic, physical and chemical properties to be used to predict the position of elements
69
atomic radii:
-> a measurement of the size of its atoms
-it is the distance from the nucleus to the boundary of an electron cloud
-a covalent radius is measured between two bonded atoms
-a van der Waals radius is meaured between two non-bonded atoms -> for Noble gases
-the metallic radius is measured between two metal atoms
70
atomic radii: across a period
-atomic radii decreases across a period
-electrons are being added to the same energy level and protons are being added to the nucleus
-this resulted in an increased nuclear charge therefore the attratcive force between the nucleus and electrons increases
-this force counterbalances the increases in electron-electron repulsion that would occur due to additional electrons in the quantum shell
71
atomic radii: down a group
-atomic radii increases down a group
-the number of occupied quantum shells increases making atoms bigger
72
melting and boiling points:
-elements with gianta lattice structures have high melting and boiling temperatures
-elements with simple molecular structures have low melting and boiling temperatures
73
ionsation energy: across a period
-as you go across a period there is an increase in the first ionisation energy
-more protons are being added to the nuclei of the atoms, this results in an increase in nuclear charge
-the electrons in the outer energy levels will be more tightly held (decrease in energy) and more difficult to remove
-the increase in nuclear charge is stronger than the increase in electron-electron repulsion (which would decrease IE)
74
ionisation energy: down a group
-as you go down a group the first ionisation energy decreases
-electron removed from the outer energy level is increasingly distant from the nucleus as a result of the additional quantum shell, the attraction of the positive nucleus diminishes and it becomes easier to remove
-the outer electrons experience increased repulsion from the inner electrons which leads to a decreased IE
-this pattern exists for groups 1-2 and 5-8
75
anomalies in IE:
-for example, the first IE of Boron is lower than Beryllium-this is unexpected
-although the nuclear charge of the boron atom is greater than that of the beryllium atom, the outer electron of boron has more energy, since it is in a 2p orbital as opposed to the 2s orbital for beryllium
-for this reason, the energy required to remove 2p electron in boron is less than the energy required to remove a 2s electron from a beryllium atom
76
cations and anions:
-most ionic compounds are formed by the direct combination of two elements:
-atoms form cations (positive ions) and anions (negative ions) by losing or gaining electrons
-for example:
2NA -> 2Na+ + 2e- (oxidation to form cations)
Cl2 + 2 e- -> 2Cl- (reductions to form anions)
77
ionic bonding:
-ionic bonding is the electrostatic attraction between the positive and negative ions
-the electrostatic interaction between ions is not directional, they simply need to be in the presence of each other to be attracted regardless of their orientation
-ionic compounds tend to be solid materials consisting of a regular arrangement of oppositely charged ions to form a giant lattice structure
-the most common one is sodium chloride
78
strength of ionic bonding:
-the size of an ion determines how well it can pack together
-the more closely packed the ions are, the stronger the electrostatic interactions and therefore the stronger the bonding
-in general, the smaller the ion and the larger the charge, the stronger the ionic bonding
79
evidence for the existence of ions: experiments
-we can use experiments such as electrolysis to prove the existence of ions
-one common experiment is passing a direct current through copper(II) chromate(VI) solution is green in colour
-Cu2+ ions are blue and migrate to the negative electrode (cathode) turning the surrounding solution blue
-CrO4 2- ions are yellow and migrate to the positive electrode (anode) turning the surrounding solution yellow
80
evidence for the existence of ions:
-electron density maps, such as the one shown below of NaCl, also show two distinct ions of different sizes
81
polarisation and polarising power:
-in an ionic lattice, the positive ion will attract the electrons of the anion
-as electrons are pulled towards the cation, the anion undergoes a distortion of its electron density causing a region where the electrons exist in an area of orbital overlap
-this distortion is known as polarisation
-polarisation is caused by:
-high charge and small size of the cation
-high charge and large size of the anion
81
ionic radii:
-down a group the ionic radii increases as the ions get more quantum shells therefore they are larger
-these are isoelectronic elements (same electron configuration)
-as the number of protons increases, the ionic radius decreases
-as the positive charge increases, the electrons are more strongly attracted and are therefore pulled closer to the nucleus
82
charge density: cation
-the ability of a cation to attract electrons from the anion towards itself is called the polarising power
-a cation with a high charge and small size has a high polarising power
-polarising power can be estimated by measuring the charge density
charge density = charge/r^2
83
charge density: anion
-the ease with which an anion is polarised depends on the size and charge
-a large anion with a high charge are polarised the most easily
-polarisation causes a slight overlap and sharing of electrons, this causes a degree of covalent bonding to exist in an ionic substance
84
physical properties of ionic substances:
-melting temperatures
-brittleness
-electrical conductivity
-solubility
85
melting temperatures of ionic substances:
-ionic substances are made of huge lattices structures of oppositely charged ions
-due to the combined forces of attraction between all these ions, the bonds in an ionic lattice are very strong
-a large amount of energy is required to overcome these bonds for the ions to break free
-this causes ionic substances to have high melting points
86
brittleness of ionic substances:
-stress can be applied to a section of the ionic crystal
-this causes the layers of ions to slide over each other which causes ions of the same charge to be side by side
-the ions repel each other and break apart
87
conductivity:
-solid ionic compounds do not conduct electricity as their ions are held in place in the lattice structure and there are no delocalised electrons
-molten ionic compounds will conduct as the ions become mobile and will migrate to the electrode with a charge opposite to their own
-aqueous solutions of ionic compounds also conduct electricity (they undergo electrolysis) and the lattice breaks down into separate ions
88
solubility of ionic substances:
-the majority of ionic compounds are soluble in water
-the energy required to break apart the lattice structure can be supplied by the hydration of the separated ions
-positive and negative ions are both attracted to the water due to its polarity
89
covalent bonding:
-a covalent bond is the electrostatic attraction between the shared pair of electrons and the nuclei of two non-metals
-he shared pair of electrons then allow each atom in the bond to have a full outer electron shell
-the way the orbitals overlap depends on their shape
-sigma bonds - end on end
-pi bonds - side by side
90
covalent bonds: types of bonds
-a pi bond can only be performed after a sigma bond is formed
-for this reason, pi bonds only exist in double or triple bonds
-in a double bonds, one bond is a pi and the other is a sigma
-in a triple bond, two bonds are pi and the other is a sigma
91
bond length and bond strength:
-the bond length is the distance between the nuclei of two covalently bonded atoms
-the strength of a covalent bond is measured as the amount of energy required to break one mole of the bond in the gaseous state
-in general, the shorter the bond length, the stronger the bond due to the increased electrosatic attractions and overlap of the orbitals
92
electronegativity:
-> is the ability of an atom to attract a bonding pair of electrons
-electronegativity increases across a period due to an increase in nuclear charge as you move across a period from left to right (i.e. it has a greater attraction for the bonding electrons)
-electronegativity decreases down a group due to the addition of another energy level of electrons as you go down a group
-these extra energy levels shield the bonded electrons from the nucleus; therefore they are not as strongly attracted
93
polar bonds:
-if electrons in a covalent bond are unequally shared, they will form a polar bond
-polar covalent bonds still have an overlap of orbitals the same as pure covalent bonds
-the difference however is that the electrons are pulled towards the more electronegative atom
-this causes a slightly negative charge on the more electronegative atom and a slightly positive charge on the less electronegative atom
94
bonding continuum:
-the two ideals of bonding are pure covalent and pure ionic and polar covalent bonds make up the spectrum in between these
-a polar covalent bond is a covalent bond which has some ionic character
-if the difference in electronegativity is large enough then the bond will be ionic as the pull for the electrons is so unequal they are transferred between atoms instead of shared
95
discrete covalent bonding:
-a discrete (simple) molecule is an electrically neutral group of two or more atoms held together by covalent bonds
-these molecules can be displayed as dot and cross diagrams
96
discrete covalent bonding: multiples
-we can also show multiple bonds (such as double or triple) using dot and cross diagrams
-these are also known as displayed formula as it shows each bonding pair in the compound as a line between the two atoms involved
97
dative covalent bonds:
-> is formed when an empty orbital of one atom overlaps with an orbital containing a non-bonding pair (lone pair) of electrons of another atom
-the bond is shown by an arrow from the atom providing the non-bonding pair to the atom with the empty orbital
-common examples: hydroxonium and ammonium ion
98
dative covalent bonds: and dimers
-another example is aluminium chloride
-this molecules has only 6 electrons in its outer shell
-just above its sublimation temperature (in the gas phase) aluminium chloride exists as Al2Cl6
-one of the atomic orbitals of a chlorine in one of the AlCl3 molecules overlaps with an empty orbital of the aluminium atom of the other molecukes
-the same happens with the chlorine of the second molecule and the aluminium of the first
-this forms two dative bonds in the dimer
-this molecules is known as a dimer
99
electron pair repulsion theory (EPR):
-EPR theory states the following:
1. the shape of the molecule or ion is caused by the repulsion between electrons (bonding and non-bonding) around the central atom
2. the electron pairs arrange themselves around the central atom so the repulsion is at a minimum (or maximum separation)
3. lone pair-lone pair repulsion > lone pair-bond pair repulsion > bond pair-bond pair repulsion
100
shapes of molecules of ions:
-linear
-trigonal planar
-tretrahedral
-trigonal pyramidal
-octahedral
-trigonal pyramidal
-v-shaped
-each shape of a molecule or ion also has a bond angle associated with it
101
shapes: linear
-bond pair: 2
-lone pair: 0
-bond angle: 180º
102
shapes: trigonal planar
-bond pair: 3
-lone pair: 0
-bond angle: 120º
103
shapes: tetrahedral
-bond pair: 4
-lone pair: 0
-bond angle: 109.5º
104
shapes: trigonal bipyramidal
-bond pair: 5
-lone pair: 0
-bond angle: 90º/120º
105
shapes: octahedral
-bond pair: 6
-lone pair: 0
-bond angle: 90º
106
shapes: trigonal pyramidal
-bond pair: 3
-lone pair: 1
-bond angle: 107º
107
shapes: v-shaped
-bond pair: 2
-lone pair: 2
-bond angle: 104º
108
polar and non-polar molecules:
-differences in electronegativity between atoms in a bond cause a drift of bonded electrons
-this separation of charge is known as a dipole
-each bond in a molecule will have a dipole associated with it
-these dipoles are known as polar bonds
-depending on the bonding angles, the individuals dipoles of each bond can either reinforce each other or cancel each other out
-if the dipoles cancel each other out, the molecules will have no overall dipole and the molecule is termed non-polar
-if the dipoles reinforce each other, the molecule will have an overall dipole and the molecule is termed polar
109
trigonal planar shape: example
-boron chloride (BCl3)
-polar bonds
-non-polar molecule
110
linear shape: example
-carbon dioxide (CO2)
-polar bonds
-non-polar molecule
111
tetrahedral shape: example
-trichloromethane (CHCl3)
-polar bonds
-polar molecule
+
-tetracgloropmethane (CCl4)
-polar bonds
-non-polar molecule
112
bent shape: example
-water (H2O)
-polar bonds
-polar molecule
113
trigonal pyramidal shape:
-ammonia (NH3)
-polar bonds
-polar molecule
114
metallic bonding:
-metals have a regular arrangement of cations surrounded by delocalised electrons
-the electrostatic forces of attraction between the cations and electrons are called metallic bonding
physical properties:
-high melting temp.
-good electrical conductivity
-good thermal conductivity
-malleability
-ductile
115
melting point:
-to melt a metal, the forces of attraction between the cations and delocalised electrons must be overcome
-this must be done to such an extent that the cations become free to move around
-metals have a lattice structure therefore large amount of energy are needed to overcome the attractions
116
melting point: group 1, 2 and d block
-the number of delocalised electrons per cation will determine the melting temperature of the metal
-group 1 metals (1 delocalised electron per cation) have low melting points
-group 2 metals (2 delocalised electron per cation) have higher melting points
-D-block elements have higher number od delocalised electrons per cation therefore even higher melting points
117
melting point: charge-to-charge radius
-melting temperatures are also affected by the charge-to-charge radius of the cation
-the greater the charge-to-radius ratio, the stronger the attraction for the delocalised electrons
-this means that for cations of the same charge, the smaller one will have a higher melting temperature
-this is why lithium has a higher temperature than sodium
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conductivity:
-electrical conductivity:
-when a voltage is applied to the metal, the delocalised electrons become attracted to the positive terminal of the cell- this is a flow of electric current
-thermal conductivity:
-the delocalised electrons pass kinetic energy along the metal so heat can move through it
-the cations are closely packed and pass kinetic energy from one cation to another
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malleability and ductility:
-metals can be hammered or pressed in to different shapes (malleability) or drawn into a wire (ductility)
-when a stress is applied, the layers of cations can slide
-the delocalised electrons move with the cations and there is repulsion between the cations in each layer
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metallic lattices:
-these are composed of a regular arrangement of cations surrounded by delocalised electrons
-they typically follow properties:
-high melting and boiling temperatures
-good electrical conductivity
-good thermal conductivity
-melleability
-ductility
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ionic lattices:
-these are composed of a regular arrangement of cations and anions
-they typically have the following properties:
-fairly high melting temperatures
-poor electrical conductivity when solid but good when molten
-brittleness
-often soluble in water
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covalent lattices:
-these are also known as network covalent lattices, they consist of giant network of atoms linked by covalent bonds
-the four most common are: diamond, graphite, graphene and silicon (IV) oxide
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diamond:
-in diamond, each carbon has 4 sigma bonds to four other carbon atoms
-it has a tetrahedral shape with bond angles of 109.5º
-it has an extremly high melting point as the strong C-C bonds require large amount of energy in order to be broken
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graphite:
-in graphite, each carbon atom has 3 sigma bonds to three other carbon atoms forming interlocking hexagonal rings
-the 4th electron is in a p-orbital and becomes delocalised
-these delocalised electrons cause graphite to be a very good conductor of electricity
-however it can only conduct electricity parallel to its layers as. the delocalised electrons cannot move between layers
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graphene:
-graphene is bonded exactly the same way as graphite
-the one major difference is that graphene is a very think sheet of carbons (one atom thick)
-it is extremely thin however it is 200 times stronger than steel
-it is an excellent thermal conductor and it can self repair any damage
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hydrocarbons:
-> are substances that are made of hydrogen and carbon only
-they can be classified as saturated and unsaturated
-saturated hydrocarbons contain as much hydrogen as possible and are made of only single bonds
-unsaturated hydrocarbons have fewer hydrogens that the maximum and will contain at least one double or triple bond
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types of hydrocarbons:
-alkanes CnH2n+2
-alkenes CnH2n
-cycloalkanes CnH2n
-cycloalkenes CnH2n-2
-each can be classified differently and have their own structures and general formula
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functional groups:
-> is an atom or group of atoms in a molecule that give the compound its distinctive and predictable properties
-e.g. -OH, -COOH, C=C
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homologous series:
-> is a family of compounds with the same functional group and differe by CH2 from the next member
-e.g.alkanes, akenes and alcohols
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properties of homologous series:
-members of a homologous series will also share the same general formula
-all alkanes will completely combust to form carbon dioxide and water
-alcohols will show a trend in boiling temperatyre- the temp. will increase as the molar mass increases
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to name organic compounds:
1) identify the longest continuous chain of carbons
2) number the carbon so the branch or functional group is on the lowest possible number
3) identify the position of, number of and type of branches
4. name the structure starting with the branches and then the length of the alkene chain
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structural isomers:
-isomers are compounds which have the same molecular formula but different structural formula
-e.g. all have to have C5H12
-however they have different structural or displayed formulae
-two types:
-chain isomerism - different carbon chains
-position isomerism- different position of functional group
-isomers of alkenes have the double bond in different places in the carbon chain, they can also have branches the same as alkanes
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addition reaction:
two reactants combine to make one product
e.g. C2H4 + H2 -> C2H6
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substitution reaction:
two reactants combine to make two products
-e.g. C2H5Br + OH- -> C2H5OH * Br-
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polymerisation:
many monomers combine to make a polymer
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homolytic fission:
-shared pair of electrons in a covalent bond are divided equally between two atoms
-forms two free radicals
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heterolytic fission:
-shared pair of electrons in a covalent bond are kept by one of the atoms (higher electronegativity)
-forms a positive and a negative ion
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electrophiles:
-> are species that are attracted to a region of high electron density
-electrophiles are positive ions and they are attracted to a region of high electron density in another molecule
-region labelled as delta -
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hazard and risks:
-> a hazard is something that can cause harm to a user
-a risk is the chance of something causing harm
-learn symbols and meanings!!
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risk assessments:
->is the identification of the hazards involved in carrying out a procedure and the control measures needed to reduce the risk of these hazards
e.g. measures: eye protection, gloves, keeping cap on a bottle, keeping substance away from heat and how to deal with spillages
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what is crude oil?
-> is a mixture of hydrocarbons with different carbon chain lengths
-the chain lengths that can vary from 1 carbon to over 100
-it is finite, non-renewable source that cannot be replaced in a human lifespan
-it is formed from the remains of ancient marine life
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properties of hydrocarbons:
-as the chain length of hydrocarbons increase, the following trend are seen
-boiling point increases
-liquid become less volaile (how easily it evaporates)
-liquids become more viscous (how easily it flows)
-liquids become darker
-flammability decreases
-these properties happen because the intermolecular forces increase as the carbon chain length increases
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fractional distillation:
-crude oil itself has no uses and has to be separated into fractions
-all fractions are mixtures of carbon chain lengths but they have similar boiling points
-the separation process is called fractional distillation
from small molecules to large molecules
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fractional distillation: process
-the fractional distillation column has a temperature gradient
-the crude oil is heated until it boils
-vapours pass into the fractioning column (cool at the top and hot at the bottom)
-depending on it's boiling point,a hdyrocarbon will travel up the column and condense for collection
-fractional distillation separates crude oil based on the boiling point of each hydrocarbon
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cracking:
-> is a process in which long chain alkanes are converted in to alkenes and shorter chain alkanes
-the hydrocarbon is heated to give a gas and passed over a catalyst of silicon dioxide (silica) and aluminium oxide (alumina) at 600-700ºC
-the hydrocarbon (usually paraffin) is soaked into mineral wool as this does not react with the catalyst and it helps keep the liquid in its place
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cracking: process
-in this experiment, the hydrocarbon and the catalyst are heated on and off
-the delivery tube must be disconnected from the apparatus before the heating is stopped to avoid suck-back
-suck-back happens because the air pressure inside the delivery tube drops and pulls the water back into it
-if this happens, the test tube can shatter due to cold water touching the hot glass
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cracking: number of carbon atoms
-the number of carbon atoms and hydrogen atoms on each side of the equation are the same
-this will form at least one unsaturated alkene because there are insufficient hydrogen atoms to form two alkanes
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reforming:
-> is the conversion of straight-chain hydrocarbons into branched-chain and cyclic hydrocarbons
-this process usually involves heating the hydrocarbons with a catalyst such as platinum
-it usually happens in an engine to help the fuel burn more smoothly
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combustion of hydrocarbons:
-a fuel is substance which, when burned, releases heat energy
-the burning of fuels is known as combustion
-there are two types of combustion: complete and incomplete
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complete combustion:
-> complete combustion of a hydrocarbon will form carbon dioxide and water
-water is not an issue as a product of combustion
-however carbon dioxide is a greenhouse gas and is responsible for climate change
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incomplete combustion:
-incomplete combustion of a hydrocarbon will form carbon monoxide (or carbon as soot) and water
-carbon monoxide is poisonous as it reduces the ability of blood to carry oxygen around the body
-tiny particles of carbon in the atmosphere are also harmful
-sometimes the fuel does not burn at all and some of it is released unchanged- known as unburned hydrocarbons (UHC)
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oxides of sulfur:
-some molecules in crude oil can contain sulfur impurities
-when these are burned they form sulfur dioxide which can then in turn react in the atmosphere to form sulfur trioxide
-these are acidic gases so they will dissolve in water to form sulforous and sulfuric acid- causing acid rain
153
oxides of nitrogen:
-very few molecules contain nitrogen impurities and when near contact with spark plugs- very high temperature
-these will form oxides of nitrogen (NOx)- the main ones are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
-nitrogen dioxide can dissolve in atmospheric water to form nitrous acid and nitric acid- causing acid rain
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catalytic convertors:
-> are attached to car exhaust systems to help reduce the pollution caused
-they typically contain small quantities of platinum, palladium or rhodium over a honeycomb mesh
-these help remove carbon monoxide, unburnt hydrocarbons and oxides of nitrogen
-sulfur is not well removed by these systems so they are removed before the fuel is burned
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alternative fuels: biofuels
-as well as pollution from the combustion of fuels, we also must consider concerns such as depletion natural resources and climate change
-we need to use alternative fuels such as biofuels as they are renewable sources
-the closer a fuel can to carbon neutral the better- carbon neutral means a zero net effect on the amount of carbon dioxide in the air
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alternative fuels: bioalcohols or hydrogen
-example of alternative fuels are bioalcohols or hydrogen
-determining the best alternative fuel requires many factors to be considered
-for biofuels, the factors are- land use, yield, manufacture, transport and carbon neutrality
-hydrogen is also useful fuel as when burned it forms water which is not harmful
-however, obstaining the hydrogen gas from water requires energy (from fossil fuels) and storage is also a problem
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substitution reactions:
-alkanes are particularly unreactive (with the exception of combustion) due to the fact they contain only single bonds
-alkanes do undergo subsitution reactions-most commonly with a halogen, this process is called halogenation
-e.g. methane reacting with chlorine:
CH4 + Cl2 -> CH3Cl + HCl
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chlorination of methane: initiation
-the chlorination of methane occurs in three steps: initiation, propagnation and termination
Step 1: initiation
-ultraviolet light breaks the chlorine molecile into two chlorine free radicals through homolytic fission
Cl2 -> Cl. + Cl.
-this is sometimes shown using half curly arrows:
Cl:Cl -> Cl. + Cl.
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chlorination of methane: propagnation
Step 2: propagnation
-chlorine free radicals collide with methane molecules and remove a hydrogen to form a methy free radical and hydrogen chloride
Cl. + CH4 -> HCl + CH3.
-the methyl free radical formed removes a chlorine atom from a chlorine molecule creating chloromethane and a chlorine free radical
CH3. + Cl2 -> CH3Cl + Cl.
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chlorination of methane: termination
-two free radicals formed in the propagnation step can collide and react to form a molecule
-there are two free radicals formed, there are three possible outcomes
Cl. + Cl. -> Cl2
Cl. + CH3. -> CH3Cl
CH3. -> CH3. -> C2H6
-the sequence of reactions is stopped as two reactive species are converted into an unreactive species
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further substitution reactions:
-this reaction has formed chloromethane- one hydrogen is replaced by one chlorine
-further subsitution can also occur to form dichloromethane, trichloromethane and tetrachloromethane:
1. dichloromethane
CH3Cl + Cl2 -> CH2Cl2 + HCl
2. trichloromethane:
CH2Cl2 + Cl2 -> CHCl3 + HCl
3. tetrachloromethane:
CHCl3 + Cl2 -> CCl4 + HCl
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alkenes:
-alkenes are unsaturated hydrocarbons containing a C=C bond
-they are generally formed in cracking reactions
-all members of the alkene homologous series share the general formula CnH2n
-due to the presence of the double bond, alkenes are much more reactive than alkanes
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naming alkenes:
1. identify the longest continuous chain of carbons
2. number the carbons so the double bond is one the lowest possible number
3. identify the position of, number of and type of branches
4. name the structure starting with the branches and then the length of the alkene chain
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C=C bonds:
-it is important to understand the two bonds in C=C are not the same-one sigma bond and one is pi bond
-the electrons in the pi bonds are further away from the carbon atoms than the sigma bonds
-this means they are more likely to undergo reactions causing the alkenes to be more reactive than alkanes (sigma bonds only)
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gemoteric isomerism:
-geomteric isomers differ from each other because their atoms or groups are attached at different positions on opposite sides of the C=C bond
-this type of isomerism only exists in alkenes as the presence of the C=C bond leads to restricted rotation
-this means the groups attached to each C=C bond can only be in one of two position
-we use two naming conventions for these: cis/trans or E/Z
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cis/trans and e/z
similar groups: cis/trans
-groups on opposite sides = trans
-groups on the same side = cis
different groups: e/z
-if imporant halogens are in below double bond = z isomer
-if important halogens are opposite double bond = e isomer
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addition reactions:
-a C=C bond is made up of sigma and pi bonds - a sigma bond is stronger than a pi bond but it is not twice as strong
-most reactions of alkenes involve the double bond becoming a single bon
-this means the sigma bond remains unchanged and the pi bond electrons are used to form bonds with an attacking molecule
-this forms a saturated product with only sigma bonds
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test for alkenes:
-the chemical test for an alkene is the bromine water test
-when this reaction occurs, the pi bond in the alkene breaks and the bromine is added across the double bond
-this causes the bromine to decolourise forming colourless products (e.g. ethen + bromine -> bromoethane)
-this is known as an addition reaction
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addition reactions: examples
-hydrogenation- addition of hydrogen to form an alkane
-halogenation- addition of a halogen to form dihalogenoalkanes
-hydration- addition of water to form an alcohol
-addition of a hydrogen halide to form halogenalkanes
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oxidation to diols:
-the oxidation of an alkene to a diol includes both addition and oxidation reactions
-a diol is a compound that contains two -OH groups
-the oxidising agent is generally potassium manganate(VII) in acidic condition (dilute sulfuric acid)
-the potassium manganate(VII) provides an oxygen and the water provide the hydrogen and 2nd oxygen
-this causes two -OH groups to be added across the double bond
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oxidation to diols: equation for formation
-the equation for the formution of a diol from ethene is:
CH2=CH2 + [O] + H2O -> CH2OH-CH2OH
-the product of this reaction is ethane-1,2-diol
-during the reaction, the potassium manganate(VII) changes from purple to colourless
-it can also be used as an identifying test for alkenes as an alkane will not react in this way
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mechanisms of addition reactions:
-the mechanism of addition reactions will contain curly arrows to show the movement of the electron pair
-the delta+ end of the HBr molecule is described as an electrophile
-this electrophile will therefore attack centres of negative charge- in this case the pi bond
-the curly arrows must either:
-start from a bond and move to an atom
-start from a lone pair of electrons and move to an atom
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electrophilic addition of hydrogen halide:
-this type of reaction is known as electrophilic addition
-the reaction begins with heterolytic fission to allow the electrons to both go to one atom
-this forms a positive carbocation intermediate and a negative bromide ion with a lone pair
-the lone pair of electrons is then donated from the bromide ion and forms a covalent bond with the carbon atom in the carbocation intermediate
-this then forms the halogenoalkane in this case bromoethane
LEARN THE STEPS PLS
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electrophylic addition: asymmetric molecules
-electrophilic additions proceed via carbocations
-carbocations can be primary, secondary or tertiary
-the stability of a carbocation is greatest for tertiary carbocations and least stable for primary carbocations
-carbocations are more stable when there are more electron-releasing alkyl groups attached to the carbon with the positive charge
-the major product formed is from the more stablee carbocation
175
addition polymerisation:
-molecules with a C=C double bond can be polymerised
-they are examples of addition reactions as the alkene molecules add together to form a long chain polymer
176
naming polymers:
-the alkene molecules in this reaction are known as monomers
-the standard way to name a polymer formed is the write poly followed by the name of the monomer in brackets
177
equations for polymerisation reactions:
-polymers do not have a fixed molecular mass so their reaction are written differently to other reactions
-we use the letter 'n' to represent the number of monomer molecules reacting and then show the repeat unit inside square brackets with a subscript 'n' and extension bonds
ethene -> poly(ethene)
with PVC:
chloroethene -> poly(chloroethene)- PVC
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identifying the monomer:
-sometimes you may be given the polymer and asked to determine the monomer
-to do this, you should identify the repeated part of the structure.
-this will be 2 carbon atoms and the 4 groups joined to them
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uses of polymers:
-polymers (or plastics) are used for many everyday objects and are increasing in use because:
-polymers can be manufactured on a large scale in complex shapes with specific properties
-plastics are lighter in weight than traditional alternatives
-they are unreactive so can be used to contain many substances safely for long periods
180
polymer waste:
-some companies are now trying to develop biodegradable polymers-sometimes known as biopolymers
-these are broken down by microbes in the environment to avoid them being wasted
-they have some disadvantages such as the need for land in order to grow the plants to make the biopolymers
-when they break down, the hydrogen and carbon atoms formed cannot be directly used again
