Module 2 Flashcards
1
Q
What are the definitions of:
RIM
RAM/Ar
RMM/RFM/Mr
A
RIM - Relative isotopic mass is mass of an isotope relative to 1/12th mass of an atom of carbon - 12 (p+n)
RAM/Ar - Weighted mean mass (of all isotopes of an element) relative to 1/12th mass of a carbon - 12 atom.
FMM/Mr: relative molecular mass
MFM: relative formula mass
2
Q
What is relative molecular mass?
A
(Mr) The weighted mean mass of a molecule of a compound compared with 1/12th of the mass of an atom. carbon - 12.
3
Q
23Na
11
What is the atomic mass no. and atomic no?
A
23 = atomic mass number number (A) 11 = atomic number (Z)
4
Q
What is the atomic no. an element?
A
The no. protons in the nucleus of an atom
5
Q
What is atomic mass measured in and how is carbon-12 used to work it out?
A
- Mass of carbon-12 isotope defined as exactly 12 atomic mass units (12u)
- Standard mass for atomic mass is 1u, mass of 1/12th of an atom of carbon-12
- 1m = mass of a proton/neutron
6
Q
What is the mass no?
A
the sum of (no. nucleons) no. protons and neutrons in the nucleus - know as NUCLEON no.
7
Q
What is an isotope?
A
Atoms of the same element. Have the same no. protons diff. no. neutrons.
Same atomic no. diff. mass no.
8
Q
State the difference, if any, inthe chemical properties of isotopes of the same element.
Explain your answer
A
the chem properties are the same -they’re determined by the no./arrangement of electrons, which is the same
9
Q
State any diff and similarities in the atomic structure of the isotopes of an element. State the diff, if any, in the chem of these isotopes. Explain
A
Diff isotopes of the same element have the same proton/atomic no. but differ int he no. neutrons (also mass no).
There’s no chem diff between isotopes of the same element as they have the same no electrons/electron configuration.
10
Q
Definitions of:
a) element
b) compound
c) atom
d) molecule
e) symbol
A
a) 1 of the 100 or so simplest substances which can’t be broken down into simpler substances.
b) A substance made by chemical combination of s/more elements.
c) the smallest particle of an element that can take part in a chemical reaction
d) the smallest particle of a compound/element that can have a separate existence and has covalent bonds
e) a letter or pair of letters which represents 1 atom of an element
11
Q
f) ion
g) formula
h) equation
A
f) an electrically charged atom/group of atoms. In chemical situations ions of 1 charge are usually balanced by ions of the opposite charge e.g. in a crystal/a solution
g) a group of symbols and usually no. which represent the no. atoms of each element in either 1 molecule of a compound/a formula unit. A giant structure. A formula can also be used to show the charge of an ion/the no. atoms and the charge. A compound ion ( 1with more than 1 atom)
h) a group of formulas and usually no. which represents the no. atoms and molecules involved in a chemical reaction.
12
Q
What is the difference between cations and anions?
A
Cations: +ve ions
atoms with fewer electrons than protons
Anions: -ve ions
Atoms with more electrons than protons
13
Q
Common ion symbols
1) hydrogen
2) sodium
3) silver
A
1) H+
2) Na+
3) Ag+
14
Q
4) potassium
5) lithium
6) ammonium
A
4) K+
5) Li+
6) NH4+
15
Q
7) barium
8) calcium
9) copper (II)
A
7) Ba2+
8) Ca2+
9) Cu2+
16
Q
10) magnesium
11) zinc
12) lead
A
10) Mg2+
11) Zn2+
12) Pb2+
17
Q
13) iron (II)
14) iron (III)
15) chloride
A
13) Fe2+
14) Fe3+
15) Al3+
18
Q
16) bromide
17) fluoride
18) iodide
A
16) Cl-
17) F-
18) I-
19
Q
19) HCO3
20) OH-
21) NO3-
A
19) hydrogencarbonate
20) hydroxide
21) nitrate
20
Q
22) O2-
23) S2-
24) SO4 2-
A
22) oxide
23) sulphide
24) sulphate
21
Q
25) CO3 2-
26) CN-
27) PO4 3-
Tetra means 4
A
25) carbonate
26) cyanide
27) phosphate
22
Q
What is mass spectra used to identify?
A
The molecular mass of an organic compound and to gain further info. about its structure.
% abundances of the isotopes in a sample of an element.
23
Q
How does a mass spectrometer work?
A
- A sample is placed in the mass spectrometer
- Sample is vapourised and then ionised to form =ve ions
- Ions are accelerated. heavier ions move more slowly and are more difficult to deflect than lighter ions, so the ions of each isotope are separated.
- ions are detected on a mass spectrum as a mass-to-charge ratio m/z. Each ion reaching the detector adds to the signal, so the greater the abundance, the larger the signal.
24
Q
mass- to - charge m/z =
A
relative mass of ion/relative charge of ion
For 1 ion with 1 +ve charge, this ratio is equivalent to the relative isotopic mass, this recorded on the x-axis of a mass spectrum. Each peak show the % abundance of each isotope.
25
Waht are the binary compounds?
They contain 2 elements only.
- To name, use the name of the 1st element but change the ending of the second element to ide.
- for ionic compounds, the metal ion always comes first.
e. g. sodium and oxygen form sodium oxide
26
What are polyatomic ions?
```
An ion containing more than one element bonded together
e.g. ammonium NH4+
nitrate NO3
carbonate CO3 2-
phosphate PO4 3-
```
27
What are the 7 diatomic molecules?
```
H2 - Have
N2 - no
F2 - fear
O2 - of
I2 - ice
Cl2 - cold
Br2 - beer!
```
28
```
Give the oxidation numer and some exaples of the following combined elements.
O
F
Mg2+, Ca2+
H
Na+, K+
Cl-, Br-, I-
```
combined oxidation no. Examples
element
O -2 H2O, CaO
H +1 NH3, H2S
F -1 HF
Na+, K+ +1 NaCl, K2O
Mg2+, Ca2+ +2 MgCl2, CaO
Cl-, Br-, I- -1 HCl, KBr, CaI2
29
Special cases: Give the oxidation no. of these:
i) H in metal halides
ii) O in peroxides
iii) O bonded toF
i) -1 e.g. NaH, CaH2
ii) -1 e.g. H2O2
iii) +2 e.g. F2O
30
What is oxidation and reduction?
```
Oxidation - is addition of oxygen.
- loss of electrons
- increase in oxidation number
Reduction: removal of oxygen.
- gain of electrons
-decrease in oxidation number
IF one happens so must the other
```
31
Redox: loss/gain of oxygen
What is happening in the reaction?
CuO(s) + H2 (g) - Cu(s) + H2O (l)
Copper (11) oxide has lost oxygen it has therefore been reduced.
Hydrogen has gained oxygen it has therefore been oxidised.
32
Redox: electrons
What is happening in the reaction?
2Fe(s) + 3Cl2(g) - 2FeCl3 (s)
FeCl3 contains +ve and -ve ions, Fe3+ and Cl-.
-Iron loses electrons and is oxidised. 2Fe - 2Fe3+ + 6e-
- Chlorine gains electrons and is reduced. 3Cl2 + 6e- - 6Cl-
The electrons gained and lost balances:
2Fe in2Fe - each Fe looses 3e- - total of 6 electrons lost
6 Cl n 3Cl2 - each Cl gains le- - total of 6 electrons gained.
33
Redox: oxidation no.
What is happening in this reaction?
Cu(s) + 2AgNO3(aq) - 2Ag(aq) + Cu(NO3) (aq)
Cu(s) + 2AgNO3(Aq) - 2Ag(aq) + Cu(NO3)2(aq)
oxidisation-
reduction
+1 0 +2
The changes in ox.no. apply to each atom and the total changes in ox. no. balance.
1Cu in Cu - Cu increase by +2 - total increase = +2
2Ag in 2AgNO3 - each Ag decreases by -1 - total decrease = -2
34
Redox: reactions of acids:
What's happening in these reactions?
A) Zn(s) + 2HCl(aq) - ZnCl2(aq) + H2(g)
B) 2Al(s) + 3H2SO4(aq) - Al2(SO4)3(aq) + 6H2(g)
Look at card for answer
35
What are shells?
- energy levels
- energy increases as shell no. increases
- shell no/energy level no. is called the
PRINCIPAL QUANTUM NUMBER, n
36
Sub Shells
sub shells max no. electrons
s 1.
p 2.
d 3.
f 4.
1. 2
2. 6
3. 10
4. 14
37
Describe how orbitals are filled up
* each at ... orbital can hold max. 2 electrons
* electrons spin clockwise/anticlockwise
* electrons can only occupy the same orbtial if they're opposite/paired spins
* filled in a certain order to produce the lowest energy arrangement possible
* filled in order of clockwise energy
* where there's more than 1 orbital with the same energy, they've first filled singly by electrons (clockwise, anti, clockwise, clockwise)
* this keeps electrons in an atom as far apart as possible
38
Why is (Fe) in 'd' block of the periodic table?
It's outermost electrons are in the 'd' orbitals OR because it's highest energy level is (3) d.
39
Ionic properties
Solubility: Ionic lattice must be broken down. H2O molecules must attract and surround the ions. Insoluble in non-polar solvents. soluble in water.
Melting point: very high
Strength@ very brittle. any dislocation leads to layers moving and similarly charged ions being next to each other - repulsion
Electrical conductivity: don't when solid - ions held strongly in lattice - conduct when molten/or in aqueous solution - ions are mobile
40
```
Covalent networks:
(Macro (or giant) molecules)
diamond, graphite and silica
a) bonding
b) melting point
c) strength
d) electrical (exception)
```
a) Many atoms are joined together in regular array by a large no. covalent bonds. Diamond - each carbon atom is joined to 4 others.
Graphite - each carbon atom is joined to 3 others.
b) Very high - structures are made up of a large no. covalent bonds, all which need to be broken if the atoms are to be separated.
c) diamond and silica (Sio2) - Hard - exists in a rigid tetrahedral structure. graphite - sofe - consists of layers which are only attracted by weak london forces. Layers can slide over each other, it's used as a lubricant and pencils
d) don't conduct - no free ions/electrons
But graphite conducts electricity
-each atom only uses 3 of its outer shell electrons for bonding to other C atoms.
-the remaining electron can move through layers allowing the conduction of electricity.
-C atoms in diamond use all 4 electrons for bonding so have no free electrons.
41
Waht is the orbital theory?
| Covalent bonding
Covalent bonds are formed when orbitals each containing 1 electron, overlap.
This forms a region in space where an electron pair can be found; new molecular orbitals are formed.
The greater the overlap the stronger the bond.
42
```
Structures containing covalent bonds:
Simple molecules
1) bonding
2) electrical
3) solubility
```
1) Atoms joined together within the molecule by covalent bonds.
2) Don't conduct elec -they've no mobile ions/electrons
3) Tend to be more soluble in organic solvents than in water; some are hydrolysed
4) LOW - intermolecular forces are weak - Vander Waals forces caused by dipoles caused by the varying position. electrons in molecules. They increase as molecules get more electrons.
Some boiling points are higher due to additional forces of attraction.
e.g.
CH4 - 161C C2H6 - 88C C3H8- 42C
43
What is average bond enthalpy?
average energy needed to break one mole of a particular bond
44
Examples of bond types:
Chemical bonds
Physical bonds
```
chemical bonds
- strong bonds
- ionic/electrovalent
- covalent
- dative covalent
- metallic
physical bonds
- weak bonds
- van der waals' forces - weakest
- dipole - dipole interaction
- hydrogen bonds - strongest
```
45
Waht are:
a) cations
b) anions
c) electron affinity?
a) positive ions - (ionisation energy)
b) negative ions - (electron affinity)
c) the energy change when 1 mole of gaseous atoms acquires 1 mole of electrons(from infinity) to form 1 mole of gaseous positive ions
X(g) + e- - X-(g)
46
Explain the formation of hydrogen bonding:
* An extension of dipole - dipole interaction giving rise to even higher b.p.
* bond between H and 3 most electr-ve elements, F, O and N are very polar.
* because of the small sizes. H, F, N and O the partial charges are concentrated in a small volume thus - a high charge density.
* Make the intermolecular attractions greater and leads to even higher bp
47
What is the structure/formation/properties of ice?
* each H2O molecule is hydrogen bonded to 4 others in a tetrahedral formation
* ice has a diamond-like structure
* its volume is larger than the liquid H2O making it
* when ice melts, the structure collapses slightly and the molecules come closer together
* they then move a little further apart as they get more energy (warmer)
* this is why H2O has a max. density at 4C and ice floats
48
Describe the formation /properties of liquid water
* intermolecular hydrogen bonding gives higher than expected b.p.
* extra attraction between molecules just below surface give high surface tension and causes the meniscus to be the shape it is
49
What is elctronegativity?
A measure of the attraction of a bonded atom for t he pair of electrons in a covalent bond.
Increases towards F in the periodic table.
50
Elctronegativity
| Non Polar bond
C-C non polar
* similar atoms have the same electronegativity
* they will both pull on the electrons to the same extent.
* the eletrons will be equally shared
51
Polar bond
* different atoms have different electronegativities
* one will pull the electron pair closer to its end
* it will be slightly more negative than average, delta minus
other will be slightly less negative, or more positive, delta +
* a dipole is induced and the bond is said to be polar
* the greater the difference is electro positive, the greater the polarity of the bond.
52
What is a polar molecule?
* some molecules are polar if they contain polar bonds.
* the molecules will be polar if they have a net dipole movement
* its a bit like balanced forces
* non-polar molecule-dipoles in bonds within the molecule cancel each other out.
* polar molecule-dipoles do not cancel each other out
* experiment - charge rod near running liquid in burette
53
Waht is the Pauling Scale?
* A scale for measuring electronegativity
* values increase across periods
* values decrease down group
* fluorine has the highest value
H
2.1
Li Be B C N O F
1.0 1.5 2.0 2.5 3.0 3.5 4.0
Na Mg Al Si P S Cl
0.9 1.2 1.5 1.8 2.1 2.5 3.0
K Br
0.8 2.8
54
What s the occurrence of permanent dipole-dipole interaction?
*between molecules containing polar bonds in addition to the basic van der Waals
* the extra attraction between dipoles means that more energy must be put in to separate molecule
* get higher b.p.than expected for a given mass.
H
55
Explain the origin of London forces:
| Instanteous dipole-induced dipole forces
* electrons in atoms/molecules are moving at high speeds in orbitals.
* It's possible for more electrons to be on one side of an atom/molecucle than the other
* this forms a dipole where 1 side is slightly negative and the other slightly positive
* a dipole is 1 atom/molecule can then induce a dipole in a neighbouring 1
See diagram on card
56
What are the results/trends and some examples of instantaneous dipole-induced dipole forces (London forces)?
Result - atoms/moleucles become attracted to each other. this makes them harder to separate and gives them higher b.p.
Trends:- the more electrons they're in an atom/ molecule the bigger the effect
Examples: layers in graphite are held together by weak London forces so it's soft.
the b.p. of noble gases increase down the group
However, some are higher than expected.
57
What is ionic bonding
and
How are giant ionic lattices formed?
ionic bonding - electrostatic attraction between and + cations and - anions ions
giant ionic lattices - are formed because ions attract ions in all directions
Each ion is surrounded by an oppositely charges ion, forming a giant ionic lattice.
58
there are 2 exceptions to these rules (doesn't say what the rules are!)
What are they?
and what is their electron configuration?
Cr - 1s2 2s2 2p6 3s2 3p6 3d5 4s1
Cu - 1s2 2s2 2p6 3s2 3p6 3d10 4s1
59
What are dative covalent/coordinate bonds?
A shared pair of electrons has been supplied by only 1 of the bonding pairs. The shared pair were originally a lone pair
see diagram on card
60
Electron repulsion help determine a molecule shape
| What are the relative repulsions between lone pairs and bonding pairs?
bonded pair/
61
What is covalent bonding?
* shared pair of electrons (non-metal and non-metal)
| * strong electrostatic attraction between shared pair of electrons and the nuclei of the bonded atoms.
62
What is the mole concept?
* An amount of substance
* 6.022 x 10 23 of something
* The Ar/Mr of a substance in grams = 6.022 x 10 23 atoms = 1 mole
* moles = mass/Mr - molar mass is measured as gmol -1
63
What is the empirical formulae?
```
The simplest ratio of atoms in an element in a compound
eg C6 H12O2 = C3 H60
1) Find no. moles
2) Find ratio
3) Put into formulae
```
64
What is molecular formula?
The number of type. atoms each element in a molecule
65
What is Avogadros' law?
Equal volumes of all gases at the same temperature and pressure contain an equal no. molecules
Avogadro constant - 6.022 x 10 23
66
Finding the formula of hydrated salts
worked e.g. 15.67g sample of a hydrate of magnesium carbonate was heated, without decomposition to form 7.58g of anhydrous magnesium carbonate.
what is the formula of the hydrate?
1) find the mass of H2O produced therefore no. moles
15.67 - 7.58 = 8.09/18 = 0.4494moles H2O
2) Find no. moles. anhydrous product
7. 58g/(84.3) - MgCo3 Mr = 0.09 moles
3) Find ratio
0. 4494/0.09 = 5 0.09/0.09 = 1 therefore MgCo3.5H2O
67
What are the steps for working out reacting masses?
1) Write balanced equation
2) What does the equation tell you about the amount in moless of the substances you're interested in
3) Change amounts in moles to masses in grams
4) Scale masses to the ones in the question
What mass of CO2 is produced when 64g of methane is burned in a plentiful supply of air?
1) CH4 (g) + 2O2 9g) - CO2 (g) + 2H2O (l)
2) 1 mole 1 mole
3) 16g 44/16g
4) 1g
64g (44x 64/16)g = 176g
68
What is molar volume?
The volume occupied by 1 mole of any gas at a particular temperature and pressure
At standard temp (0C) and pressure (atmostphere) (STP) molar V = 22.4dm3
25C and 1 atmosphere molar V = 24dm3
Allows you to calculate the amount n moles from the V of a gas and vice versa . Assume molar v of gas at room temperature and pressure = 24dm3 mol-1
69
What is Boyle's law?
at a constant temp, the volume of a fixed mass of an ideal gas is inversely proportional to its pressure
Therefore P is directly proportional to 1/V
P= k/V
(Therefore PV=k)
where k is a constant
70
What's Charles's law?
assuming a fixed mass and a constant pressure, the volume of a gas is directly proportional to the temp (in kelvin, k)
V is directly proportional to T
Therefore V = kT
71
What s Gay-Lussac's Law?
Assuming a constant volume and mass, the pressure exhibited by an ideal gas is directly proportional to the temperature.
P is directly proportional to T
Therefore P = kT
72
What is the overall gas law/equation?
```
PV = nRT
P is pressure (Pa)
n is no. of moles
T is temp - k
R is gas constants
V is volume (m3)
8.31 Jk-1 mol -1
xcm3 = x x 10-6 m3
the ideal gas equation!
```
73
e.g. 2Na + 2H2O - 2NaOH + H2
| What volume of H2 will thee addition of 1 gram of Na produce?
```
Na
moles = 1g/23
= 0.043478 - (2:1) 0.021739 moles of hydrogen
x 24 = 0.5217391304dm3
= 521.74cm3
```
74
How to work out reacting volumes?
e.g. HCl (aq) + NaOH (aq) - H2O (l) + NaCl (aq)
50cm3 of HCl iss fully neutralised by 27.5cm3 of 0.1 mol dm--3 NaOH. What is the con. HCl?
- c.f. reacting masses. - calcu. the V/conc of a reagent based upon the molarity + V. any other reagent.
Step 1 : moles of what we have
27.5cm3 therefore n = CV = 0.1 x 0.0275
0.1 moldm-3 = 0.00275 mol
Step 2: stoicheometry: ratio
NaOH: HCl = 1:1
therefore 0.00275 mol. HCl
Step 3: find the conc. HCl
c = n/v = 0.00275/0.05 = 0.055 moldm-3 HCl
75
What is the ideal gas?
* Gas particles have no volume (point masses)
* No forces. attraction exist between gas particles, except when they collide.
* At room temp and pressure (RTP) 1 mole of gas occupies 24 dm3
76
Equivalent amounts:
1 dm3 = ? cm3 = ?l = ?ml
1 cm3 = ? ml
1 dm3 = 1000cm3 = 1 litre = 1000ml
1cm3 = 1ml
77
Steps to work out molar volume:
| e.g. Calculate the volume occupied by 4.4g of CO2 at room temperature and pressure
```
1) calculate the amount in moles from the mass
from (CO2) = 44
Amount CO2 = 4.4/44g mol-1 = 0.1 mol
2) Calculate the volume
1 mol CO2 at rtp has v.24dm3
0.1 mol CO2 at rtp has 2.4dm3
```
78
Working out concentrations:
80g NaOH in 1dm3 solution has concentration. 80g dm-3
concentration = concentration (g dm-3)/
(mol dm-3) molar mass (gmol-1)
amount/mol = (conc. solution/mol dm-3) x vol solution dm3
79
What is stoicheometry?
equivalents in the reaction (big numbers)
e.g.
CH4 + 2O2 - 2H2O + CO2
80
Steps for reacting masses:
| e.g. 100g methane is burned in excess O2. How much CO2 is produced?
1) Calculate no. moles of CH4
n = m/mr = 100/16 = 6.25 mol CH4
2) Use stoicheometric ratios to find no. mol of CO2
ration = 1:1 therefore 6.25 mol CO2
3) Calculate mass of CO2
m = Mr x n
= 44 x 6.25
= 275g CO2
81
What is the percentage yield equation?
% yield = actual yield (g)/ theoretical yield (g) x 100
theoretical yield is from reacting mass calculations
82
What is atom economy?
e.g. HCl _ NaOH - NaCl + H2O
NaCl is the desired product. Calculate the reactions atom economy.
A measure of how efficient a reaction will be with regard to atoms used.
Atom economy = total mass atoms in desired product/ total mass atoms in all reactants
Mr. NaCl = 58.5 = 58.5/(40.0 + 36.5)
Mr NaOH = 40.0 = 0.76
Mr HCl = 36.5 = 76%
83
What is the method to find standard solutions?
Standard solution - conc. is known exactly
Method: solid is weighted +/- 0.01g
* solid is dissolved in a beaker using less distilled water than will be needed to fill the volumetric flask
* Put solution in volumetric flask
* Rinse trace of solution off apparatus into flask
* fill flask to graduation line with distilled water
* shake flask
84
What is water of crystallisation?
During crystalisation, water molecules can become trapped within the crystal structure. This water of crystalisation. e.g. CuSO4. 5H2O
85
How relative molecular/formula mass worked out?
RMM: Mr
compares mass of molecule with mass of an atom of C-12 and up all the RAM making up the molecule.
RFM:
compares the mass of a formula unit with mass of an atom of C-12.
+ up all the RAM of the elements in the empirical formula.
Molecular formula:
No. atoms of each element in a molecule
86
What is a standard solution?
A solution of known concentration!
| They can be prepared by dissolving an exact mass of a solute in a solvent and making up the solution to an exact volume.
87
What is the limiting reagent?
The reactant that is NOT in excess and will be completely used up first and stop the reaction.
88
What is a weak/strong acid?
A strong acid (e.g. HCl) releases all its hydrogen atoms into a solution as H+ ions and completely dissociates in aqueous solution.
A weak acid (e.g. ethanoic) only releases a small proportion of its available hydrogen atoms into solution as H+ ions and partially dissociates in aqueous solution.
89
What is a base and an alkali?
Metal oxides, metal hydroxides, metal carbonates and ammonia are all bases.
A base neutralises an acid to form a salt. An alkali is a base that dissolves in water releasing hydroxide ions (OH-) into the solution.
90
Name some common
| Acids + bases:
Acids: Bases:
H2SO4 sulphuric acid NaOH - sodium hydroxide
HCl - hydrochloric acid KOH - potassium hydroxide
HNO3 - nitric acid LiOH - lithium hydroxide
HF - hydrofluoric acid NH3 - ammonia
CH3COOH - ethanoic acid N(C2H5)3 triethylamine
H3PO4 - phosphoric acid
91
What are the Bronsted Lowry definitions of acids adn bases?
A Bronsted acid is a proton donor (H+ ions).
e.g. HCl (aq) - H+ (aq) + Cl- (aq)
A Bronsted base is a proton acceptor.
e.g. OH- (aq) + H+ (aq) - H2O (l)
92
Describe some characteristics of Acids and Bases:
Acids: Bases
* PH below 7 *PH above 7
* e.g. vinegar (ethanoic acid) * e.g. CuO
citrus fruits (citric acid) CaCO3
apples (malic acid) * H+ acceptors
* H+ donors * Alkalis produce OH- ions
in solution
* universal indicator e.g. oven cleaner
93
concentration vs. strength
| What's the difference?
Conc. acid relates too conc. H+ ions in solution
Strength - an acid determined by how readily an acid will dissociate i.e. strong acid fully disociates, weak acids don't.
e.g. HCl (aq) - H+(aq) + Cl- (aq) strong
CH3OOH interchangeable H+ (aq) + CH3OO-(aq) weak
94
What is a neutralisation reaction?
| acid + base
In neutralisation of an acid H+ (aq) ions react with a base to form a salt and neutral water.
Acid + alkali - salt + water
Acid + metal oxide/hydroxide - salt + water
Acid + metal - salt + hydrogen
Acid + metal carbonate - slat + water + carbon dioxide (g)
95
What is a titration?
A technique used to accurately measure the volume of one solution that reacts exactly with another solution.
Can be used for:
- finding conc. of a solution
- identification of unknown chemicals
- finding purity of a substance (important for quality control)
96
State the tolerance of the apparatus below:
a) 100cm3 volumetric flask
b) 250cm3 volumetric flask
c) 10cm3 pipette
d) 25cm3 pipette
e) 50cm3 burette
a) +/- 0.20cm3
b) +/- 0.30cm3
c) +/- 0.04cm3
d) +/- 0.06cm3
e) +/- 0.10cm3
97
How do you do titration calculations?
From the results you probably know both the conc. of C1 and reacting volume V1 of one of the solutions. And only the reacting volume of the other solution.
step 1: work out amount, in mol, of solute in solution for which you know both the conc. c1 and the volume v1.
Step 2: Use equation to work out the amount, in mol, of the solute in the other solution.
Step 3: Workout the unknown information about the solute in the other solution.
98
How do you carry out an Acid-base titration?
1. Add measured V of 1 solution to conical flask using pipette.
2. Add other solution to burette, record initial burette reading to 0.05cm3.
3. Add few drops of indicator to solution in conical flask.
4. Run solution in burette into conical flask sol. swirling flask to mix. Finally indicator changes colour at end point of titration. End point is used to indicate the volume of 1 sol. that exactly reacts with V of 2nd sol.
5. Record final burette reading. V. sol added from the burette = the titre = initial - final burette reading.
6. trial titration carried out first to give an approximate result.
7. Then repeated till 2 concordant (with 0.10cm3 of each other) results are recorded.
Trial 1 2 3
final burette reading/cm3
initial burette reading/cm3
titre/cm3
mean titre/cm3
