Transcript Slide 1

Reaction Rates
• During the course of a chemical reaction, reactants are
being converted into products.
• Measurement of the rate of reaction involves measuring
the ‘change in the amount’ of a reactant or product in a
certain time.
• The rate of reaction changes as it progresses, being
relatively fast at the start and slowing towards the end.
• What is being measured is the average rate over the time
interval chosen.
• Reactions can be followed by measuring changes in
concentration, mass and volume.
Where property = mass/volume/concentration
The above is used when there is no change in
mass/volume/concentration measured, for example
during a colour change reaction.
Time (s)
0
5
10
Volume (cm3)
0.00
7.00
10.50
15
20
25
12.00
12.75
13.00
30
13.00
Calculate the reaction rate;
•
•
•
•
During the first 5 seconds (7-0/5-0 = 7/5 = 1.4cm3/s)
(12-10.5/15-10 = 1.5/5 = 0.3cm3/s)
Between 10-15 seconds
Between 20-30 seconds (13-12.75/30-20 = 0.25/10 = 0.025cm3/s)
(13-0/30-0 = 13/30 = 0.43cm3/s)
Of the whole reaction
Collision Theory
A chemical reaction can only occur if there is a
successful collision between reactant molecules.
From S3 we know that we can speed up a
chemical reaction by;
1.
2.
3.
4.
Decreasing particle size (increasing surface area)
Increasing concentration (of reactant)
Increasing temperature
Adding a catalyst
Collision Theory – Particle Size
• The smaller the particle size, the higher the surface
area.
• The higher the surface area, the greater the number of
collisions that can occur at any one time.
• The greater the number of collisions, the faster the
reaction.
• Therefore the smaller the particle size, the faster the
reaction rate.
Collision Theory – Concentration
• The higher the concentration, the higher the number
of particles.
• The higher the number of particles, the greater the
chance of collisions that can occur.
• The greater the number of collisions, the faster the
reaction.
• Therefore the higher the concentration, the faster the
reaction rate.
Collision Theory – Temperature
• The higher the temperature, the higher the energy the
particles have.
• The higher the energy, the faster the particles move.
• The faster the particles move, the greater the chance
that they can collide
• The greater the number of collisions, the faster the
reaction.
• Therefore the higher the temperature, the faster the
reaction rate.
Catalysts
• A catalyst speeds up a chemical reaction without
getting used up or changed itself.
• Catalytic converters are used in exhaust systems to
turn harmful gases into less harmful gases. Platinum,
rhodium (expensive transition metals) are used as the
catalyst used in catalytic converters.
https://www.youtube.com/watch?v=rmtFp-SV0tY
The atom is made
up of a dense
centre called the
nucleus, which
contains protons
and neutrons
Modern Day Model
of the atom
The electrons are
very light and are
found in a space
around the nucleus
called the electron
shell.
Particle
Mass
Charge
Proton
1
+ve
(positive)
Neutron
1
0
(no charge)
Electron
zero
-ve
(1/1850)
(negative)
Location
16
8
Important
• Atomic Number = No. of ____________
• No. of Protons = No. of ____________
(when the atom is neutral!)
• Mass Number =
No. of __________ + No. of _________
• No. of Neutrons =
Mass Number - ____________ __________
Symbol
No. Of
Protons
No. Of
Neutrons
Atomic
Number
Mass
Number
24
12
Mg
12
45
80
35
Br
40
19
9
19
19
40
9
9
19
K
F
No. Of
Electrons
Symbol
16
8
O
31
15
P
24
12
S
No. Of
Protons
No. Of
Neutrons
Atomic
Number
Mass
Number
No. Of
Electrons
What if the element isn’t neutral?
Symbol
24
12
Mg 2+
80
Br
35
40
19
K+
16
8
O
2-
No. Of No. Of
Atomic
Mass
No. Of
Protons Neutrons Number Number Electrons
Isotopes
Isotopes
28
14 Si
25
14
31
14
Si
Si
63
29
Cu
65
29
Cu
Protons
Electrons
Neutrons
Isotopes are atoms of the same element
(same number of protons) but have
different number of neutrons.
This means for isotopes, the atomic number
stays the same but the mass number
changes.
Many elements exist as 2 or more isotopes.
The relative atomic mass (R.A.M.) of an
element is the average mass number for
a sample of that element.
R.A.M is related to isotopes.
The relative atomic mass of Copper (Cu) is 63.5
What does this tell you about the proportion of the
two types of isotopes in a sample of copper?
Electron Arrangements
Electrons are arranged in shells (or energy levels.)
Lithium has the electron arrangement 2,1 so there
are two electrons in the electron shell closest to
the nucleus, and one in the next shell:
Li
X
X
X
The 2,8,8,2 Rule
• For the first 20 elements (hydrogen to
calcium) we follow the 2, 8, 8, 2 rule.
• A maximum of 2 electrons are allowed in
the 1st electron shell, 8 electrons in the
second, 8 electrons in the 3rd and 2
electrons in the 4th.
• Each electron shell must be full before the
next one is started.
Coincidence? Maybe. Maybe Not?
The number of ______ ________ affect
the way that the atom reacts.
In other words…
Elements with the same number of outer
electrons (elements in the same group) have
similar chemical properties.
Important groups to remember;
• Group 1 – the alkali metals; all very reactive soft metals. (two
examples are ________ and ________)
• Group 2 – alkaline earth metals; similar to group 1 metals but
not as soft or reactive. (two examples are ________ and
________)
• Group 7 – the halogens; very reactive non-metals. (two examples
are ________ and ________)
• Group 8 – the noble gases; very unreactive non-metals (two
examples are ________ and ______)
• In-between groups 2/3 – transition metals. (two examples are
________ and ________)
Remember to label
what each colour
represents.
Colour in all of the solid elements one colour
(easiest thing to do is leave them white!)
Colour in all of the liquids one colour
(Bromine and Mercury)
Colour in all of the gases one colour.
(elements 1, 7, 8, 9, 17 and all of group 0)
For each of the elements (right)
write down the following;
Name –
Symbol Atomic number –
Mass number Number of protons –
Number of electrons –
Number of neutrons –
Electron Arrangement –
Group number –
Metal or non metal –
Solid, liquid or gas -
a)
b)
c)
d)
e)
f)
g)
h)
i)
j)
Carbon
Magnesium
Potassium
Fluorine
Oxygen
Nitrogen
Beryllium
Aluminium
Neon
Bromine
Why do atoms bond?
Noble gases have a complete outer electron shell.
This arrangement of electrons is very stable and
therefore other elements want to be like the noble
gases.
The number of electrons the element needs to lose or
gain to be like the noble gases is called the valency.
The atoms of other elements can collide together and
combine to achieve the full outer electron shell.
The Covalent Bond
Two atoms get close enough to
each other to collide.
O
X
X
X
X
Oxygen with red
crosses – 2,6
Oxygen with purple
crosses – 2,6
The Covalent Bond
The two atoms are attracted
to one another through the positive nucleus
of one and the negative electrons of the
other.
O
X
X
X
X
The Covalent Bond
The two atoms combine and
share enough outer electrons for each of
them to become stable (full outer shell)
X
X
X
O
Oxzygen with
red crosses –
2,8
X
X
X
Oxygen with
purple crosses –
2,8
A covalent bond is a shared pair of electrons
between non-metal atoms.
A covalent bond is held by the attraction of
the positive nucleus and negative outer
electrons of the different atoms.
Naming Covalent Compounds
If a compound name ends in ‘-ide’ then that
compound only contains two elements – e.g. carbon
nitride contains carbon and nitrogen only.
Sometimes prefixes are used in naming compounds –
e.g. silicon dioxide
Mono –
Di –
Tri –
Tetra -
one
two
three
four
Draw the following covalent elements and compounds
using both the lines and the overlapping electron
shells (circles.)
You only have to draw the outermost shell electrons.
Hydrogen Chloride (HCl)
Phosphorus Trichloride (PCl3)
Water (hydrogen oxide) (H2O)
Sulphur Fluoride (SF2)
Ammonia (nitrogen trihydride) (NH3)
Carbon Dioxide (CO2) *** tricky
1.Hydrogen Sulphide
2.Hydrogen Chloride
3.Phosphorus Oxide
4.Carbon Sulphide
5.Hydrogen Fluoride
6.Carbon Chloride
7.Silicon Oxide
8.Carbon Hydride
9.Nitrogen Hydride
10.Carbon Nitride
Shapes of Molecules
Formula = CH4
Name = Carbon Tetrahydride;
Structure is drawn like
NH3
H2O
HCl
Challenge
Your challenge is to;
a) Work out the chemical
formula
b) Name the compound (use
prefixes if necessary)
c) Draw the line drawing +
d) Draw these molecules using
the overlapping circles
methods
e) Identify the shape of the
molecule
1.Selenium and Iodine
2.Hydrogen and Fluorine
3.Carbon and Chlorine
4.Carbon and Hydrogen
5.Nitrogen and Hydrogen
6.Phosphorus and Bromine
7.Carbon and Sulphur
8.Silicon and Oxygen
Properties of Covalent Molecules
Methane molecule (CH4)
This is a covalent molecule. Other examples include
water (H2O), oxygen (O2) and candle wax (C8H18).
Covalent molecules;
– have low melting and boiling points
– can be solid, liquid or gas at room temperature.
– never conduct electricity (in any state)
Covalent Network Substances
Some covalent substances do not have individual
molecules.
Diamond is an example of a covalent network
structure. Sand (silicon dioxide) is another example.
This is a diamond structure (carbon atoms only.)
Properties of Covalent Network
Covalent networks;
• have extremely high melting and boiling
points.
• are always solid at room temperature.
• never conduct electricity
 graphite is the exception to this rule as
although it is a covalent network it will
conduct electricity.
• Although the covalent bonds within covalent
molecules are strong, the force of attraction
between the molecules are weak.
• These weak forces of attraction don’t require a
lot of energy to break and therefore covalent
molecules have low melting/boiling pts.
• All bonds in a covalent networks are very strong
covalent bonds and it takes a lot of energy to
break these bonds – i.e. very high melting/boiling
points.
Type of covalent
substances
Melting/Boiling
Pts
Type of bond
broken
Strength of
bonds broken
Covalent
Molecular
Covalent
Network
Typical Past Paper Questions - Bonding
Which line in the table shows the properties
of a covalent molecular compound?
Which of the diagrams (left)
show the structure of a
diatomic compound?
A
B
C
D
Which diagram (left) shows
the structure of a covalent
network structure?
A
B
C
D
Ionic Bonding

When a __________ atom wants to become stable
and bond with a non-metal atom, it gives away or
loses electrons.

Metals lose electrons to get a full outer shell.

These ‘lost’ electrons don’t just disappear – instead;

Non metals gain electrons to get a full outer shell.
ion
ion
An ionic bond is between a metal and non metal. An ionic
bond is held together via the attraction between the
positive metal ion and the negative non metal ion.
http://www.youtube.com/watch?v=QqjcCvzWwww
•
Ionic compounds don’t form molecules like covalent
substances. They are arranged in large lattice structures.
•
The ________ metals ions are attracted to the _______
non metals ions.
1.
This attraction is very strong making an ionic lattice a
strong stable substance – because of this ionic compounds
are always solid at room temperature.
2. Ionic compound conduct electricity when molten (melted)
and when in solution but NOT as a solid. This is because
the ions are free to move.
3. Ionic compounds have high melting points and boiling points
normally in the range; 400oC to 1400oC.
For each of the following elements;
1.
2.
3.
4.
Draw the electron arrangement of the atom.
Draw the electron arrangement of the ion.
Work out the charge of the ion.
Name the noble gas that the ion similar to.
a) Aluminium
b) Magnesium
c) Lithium
d) Calcium
e) Potassium
f) Oxygen
g) Chlorine
h) Nitrogen
i) Sulphur
j) Bromine
Example; Beryllium
The Beryllium ion (charge ____) has the same electron
arrangement as the noble gas _________________.
The ___________ ion (charge ____) has the same electron
arrangement as the noble gas _________________.
Working out formulae of Ionic
Compounds
Complex ions are ions which contain more than one element.
We can identify complex ions from their names as they end
in ‘ite’ or ‘ate’. (the exceptions to this rule are hydroxide
and ammonium)
Na
(SO42-)
1
2
2
1
Na2 (SO42-)
Work out the chemical formula for the following ionic compounds…
Remember the roman numerals indicate the valency of transition
metals.
(Be careful some contain complex ions – see data booklet)
oSodium
Lithium
Chloride
Carbonate
oCalcium
Aluminium
Oxide
Sulphate
oBeryllium
Beryllium
Bromide
Fluoride
oPalladium
Calcium (II)
Carbide
Bromide
oNickel
Iron (III)
(II) Oxide
Iodide
oTitanium
Copper (II)
(III)Sulphite
Iodide
oCaesium
Zinc (I)
Sulphite
Selenide
oBarium
Calcium
Hydroxide
Chromate
oPotassium
Sodium Carbonate
Phosphate
oSilver
Lithium
Nitrate
(I) Oxide
oIron
Hydrogen
Sulphate
(III) Phosphide
Manganese (III) Chloride
Copper (III) Carbonate
Rhodium (I) Oxide
Vanadium (II) Sulphate
Barium Carbonate
Cadmium (II) Hydroxide
Ammonium Phosphate
Ammonium Dichromate
Zinc (II) Sulphide
Hafnium (I) Permanganate
Aluminium Ethanoate
Potassium iodide reacts with lead (II) nitrate to
form a yellow solid and a clear solution. The
products are thought to be lead (II) iodide and
potassium nitrate.
Word equation
Chemical equation
Molten iron is used to join steel railway lines
together.
Molten iron is produced when aluminium reacts
with iron (III) oxide. Another product is thought
to be aluminium oxide.
Word Equation
Chemical Equation
Calcium carbonate reacts with hydrogen chloride to
form a calcium chloride, water and a gas. The gas
was tested and is turned limewater from colourless
to milky.
Write the chemical equation for the above
reaction.
Balancing Equation Examples
Not only do we always have to have the same elements on
both sides of a ‘reaction arrow’ but we also need to have the
same amounts of them too.
In order for this to happen we need to balance the
equation.
H =2 4
H=2
4
O=2
O=1
2
1.
2.
3.
4.
5.
6.
7.
Balancing Equation Practice
H2 + Cl2
HCl
Al + Cl2
AlCl3
C3H8 + O2
CO2 + H2O
Fe2O3 + CO
Fe + CO2 ***
NaOH + H2SO4
Na2SO4 + H2O
NH3 + O2
NO + H2O ***
Mg(OH)2 + HCl
MgCl2 + H2O ***
*** = tricky
Now try the sheet – do not write on it.
Formula Mass
The formula mass (or gram formula mass - gfm) of a
substance is obtained by adding the relative atomic
masses of all of the elements in a compound together.
In other words, the formula mass is the total mass
of a compound.
The formula mass has NO units.
Worked Example 1
Calculate the formula mass of calcium chloride.
Formula;
Formula Mass;
Calcium =
Chlorine =
(Total) =
Worked Example 2
Calculate the formula mass of hydrogen sulphite.
Formula;
Formula Mass;
Hydrogen;
Sulphur;
Oxygen;
(total)
Worked Example 3
Calculate the formula mass of magnesium nitrate.
Formula;
Formula Mass;
Magnesium =
Nitrogen =
Oxygen =
(Total) =
Hydrogen oxide
Calcium carbonate
Carbon dioxide
Lithium phosphate
Nitrogen fluoride
Hydrogen nitrate
Aluminium phosphide
Ammonium sulphide
Magnesium bromide
Gold (I) sulphate
Copper (II) iodide
Silver (I) ethanoate
Zinc (III) oxide
Potassium dichromate
Iron (II) chloride
Nickel (III) chromate
Lead (II) nitride
Tin (II) oxide
Mercury (I) bromide
Platinum (II) sulphite
The Mole
One mole of a substance is the formula mass
but with the units grams.
For example;
One mole of calcium chloride = 111g
One mole of hydrogen sulphite =
One mole of magnesium nitrate =
m = mass of substance
n = no. of moles
fm = formula mass
mass =
no. of mole x formula mass
m = n X fm
formula mass =
mass / no of moles
fm = m
n
no of mole =
mass / formula mass
n=m
fm
Calculate the number of moles in
10.1g of potassium nitrate (KNO3)
0.75moles of a compound X weighs 14g.
Calculate the formula mass of the substance.
What mass of sodium carbonate
(Na2CO3) is present in 0.5 moles?
Questions
1. Calculate the mass of 3 moles of copper (I) bromide. CuBr
2. Calculate the mass of 4 moles of calcium nitrate. Ca(NO3)2
3. How many moles are there in 4.23g of magnesium
carbonate? MgCO3
4. How many moles are there in 1kg of aluminium chloride?
AlCl3
5. 0.75 moles of compound Z weigh 102.35g, calculate the
formula mass.
6. 1.25 moles of compound Y weigh 69.40g, calculate the
formula mass.
How many moles of each substance in;
A – 14g of Nitrogen gas (N2)
B – 84.5g of Magnesium carbonate
C – 400g of Copper (II) oxide
D – 321g of Iron (III) hydroxide
What is the mass of;
A – 1 mole of aluminium
B – 2.5 Moles of Oxygen gas (O2)
C – 0.5 moles of Lithium sulphate
D – 0.1 moles of ethane (C2H6)
Calculation using Balanced Equation
Balanced equations can be used to calculate masses
of substances involved in chemical reactions.
When we write the balanced chemical equation, the
number in front of each formula represents the
number of moles of the substance.
e.g.
Ca +
2HNO3
Ca (NO3)2 + H2
What mass of calcium chloride would be produced when 10g of
calcium reacts fully with hydrogen chloride?
Ca
+
2HCl
CaCl2
+
H2
CH4
+
2O2
CO2
+ 2H2O
What mass of carbon dioxide is formed when 64g of methane
are burned completely in air?
Calculate the mass of iron that would be produced
from 2 moles of iron (lll) oxide.
Fe2O3
+
3H 2
2Fe
+
3H2O
C9H2O
+
14 O2
9 CO2
+ 10 H2O
Calculate the mass of water produced when 6.4g of
nonane (C9H2O) is burned.
The pH scale…
What does it really mean?
•
•
•
•
•
•
•
Red
Orange
Yellow
Green
Green/Blue
Blue
Purple
–
-
very acidic
slightly acidic
very slightly acidic
neutral
very slightly alkaline
slightly alkaline
very alkaline
Oxide
Carbon
Dioxide
Sulphur
trioxide
Sulphur
dioxide
Nitrogen
Dioxide
Acid
formed
Other
name of
acid
Carbonic Hydrogen
Acid
Carbonate
Sulphuric Hydrogen
Acid
Sulphate
Sulphurous Hydrogen
Acid
Sulphite
Nitric
Acid
Hydrochloric
Acid
Hydrogen
Chloride
Ethanoic
Acid
Hydrogen
Ethanoate
Acid
formula
Ions in
the acid
What conclusion can you come to..?
An acid is a solution which has a greater
_____________________________
______________ than pure water.
i.e.
Acids have an excess of hydrogen ions
(H+ ions)
Oxide
Name of alkali
formed
Sodium
Oxide
Sodium
Hydroxide
Potassium
Oxide
Calcium
Oxide
Ammonia
(NH3)
Ammonium
Hydroxide
Formula of
alkali
Ions in the
alkali
What conclusion can you come to..?
An alkali is a solution which has a greater
______________________________
than pure water.
i.e.
Alkalis have an excess of hydroxide ions
(OH- ions)
Neutral Substances
• Water has equal amounts of H+ and OH- ions so
the overall pH is neutral (pH 7)
• During neutralisation reactions the H+ ion (from
the acid) and the OH- ion (from the alkali) react
and produce water.
• The H+ and OH- ions being able to move is the
reason why water conducts electricity.
• All solutions contain both H+ and OH- ions.
• Acids have more H+ than OH- ions
• Alkalis have more OH- than H+ ions
• Neutral substances have equal amounts of
H+ and OH- ions
Diluting Acids or Alkalis
Diluting an acid or alkali is similar to what
happens with diluting juice.
As an acid is diluted it becomes ‘weaker’ as there
are less H+ ions. As a result the pH increases
towards pH 7 until it eventually becomes neutral.
As an alkali is diluted it becomes ‘weaker’ as
there are less OH- ions. This means the pH
decreases towards pH 7 (neutral)
Neutralisation Reactions
Acids can be neutralised using bases.
ACID + Metal Hydroxide
SALT + WATER
Worked Example
Sodium Hydroxide + Hydrochloric Acid
Sodium Chloride + Water
Bases
Bases are another word for alkalis.
Examples of bases are;
– metal hydroxides,
– metal oxides and
– metal carbonates.
Bases neutralise acids to form water.
Salts
Salts are _______ _______ compounds
(contain a positive and negative charge).
We can name the salt produced by looking at
the acid and base that have been used.
The first part of the salt’s name comes from
the base used;
e.g. Sodium Hydroxide gives – sodium …….
Calcium Oxide gives – calcium …….
Ammonia gives – ammonium…….
The second part of the name comes from the type of
acid used during the neutralisation;
Hydrochloric acid produces .......... chloride salts
Sulphuric acid produces ................. sulphate salts
Nitric acid produces ..................... nitrate salts
Ethanoic acid produces ……………… ethanoate salts
1. Complete the names of the salts produced;
a) Calcium Hydroxide + Hydrochloric Acid
b) Sodium oxide + Nitric Acid
c) Potassium carbonate + Sulphuric Acid
2. Which acid has been used to make calcium sulphate?
3. Write the word equation for magnesium hydroxide
reacting with an acid of your choice.
Metal Oxide + Acid reaction
ACID + METAL OXIDE
SALT + WATER
Example (from your experiment)
Word Equation
Chemical Equation (balance if necessary)
1.
Add 10cm3 of acid and 10cm3
of water to a 100cm3 beaker.
2.
Heat the acid until ALMOST
boiling using a slightly blue
flame.
3.
When the acid is hot enough,
use a spatula to add small
amounts of copper(II) oxide
(1g in total) to the beaker.
4.
Stir the mixture gently for up
to half a minute after each
addition.
5.
When all the copper(II) oxide
has been added, continue to
heat gently for 1 to 2min to
ensure reaction is finished.
6.
Allow the beaker to cool while
you set up the filtration.
Metal Carbonate + Acid reactions
Calcium carbonate (limestone) is often used
for building materials.
Metal carbonates react with acid, so acid rain
can damage limestone buildings.
ACID + METAL CARBONATE
SALT + WATER + CARBON
DIOXIDE
Ionic Formula
Writing ionic formula is simply writing the chemical
formula of a compound but showing the charges of
the ions.
Metals form positive/negative ions
Non metals form positive/negative ions.
Chemical Formula – NaOH
Ionic Formula – Na+ OH-
Spectator Ions
When acids and alkalis are added to water, they
dissociate to form ions.
Some of these ions react, and some do not. The
ions that don’t react are called spectator ions.
Ionic equations make it easier to identify spectator
ions. Rather than write the chemical formula of
reactants and products, we can write the ions
present in aqueous solution:
Sodium Hydroxide + Hydrochloric Acid Worked Example
Word Equation;
Formula Equation (with state symbols);
Ionic Equation (with state symbols);
Spectator ions appear on reactant and product sides of the
reaction and remain unchanged, so we can cross them out.
Rewrite equation without (omitting) spectator ions
Calculate the mass, in grams, of sulphuric acid
(H2SO4) present in 500cm3 of a 5 mol/l solution
Titration Calculation
Step 1: Write the formula:
(C x V x P)acid = (C x V x P)alkali
C = Concentration of acid or alkali
V = Volume of acid or alkali – Always make sure the
units of volume are the same on each side!!
P = Number of H+ (for acid)
Number of OH- (for alkali)
Step 2:
Find out the number of H+ in the formula of
acid. (e.g. HCl= 1 , H2SO4= 2) OR
Find out the number of OH- in the formula
of alkali. (e.g. NaOH= 1, Mg(OH)2= 2)
Step 3: Put all of the known values into the
equation shown above.
Step 4: Complete the calculation.
Examples of Power
Work out the powers of either the H+ (acids) or OH(alkalis) in the following;
•
•
•
•
•
•
•
NaOH
Ca(OH)2
H3PO4
HCl
H2SO4
Al(OH)3
H5S2
Worked Example
In a titration, 10 cm3 of 2 mol/l sodium hydroxide (NaOH)
solution was neutralised by 25cm3 of dilute hydrochloric
acid (HCl).
Calculate the concentration of the acid in mol/l.
(C x V x P)acid = (C x V x P)alkali
In a titration, 25 cm3 of 2 mol/l sodium hydroxide (NaOH)
solution was neutralised by 28.7cm3 of sulphuric acid (H2SO4).
Calculate the concentration of the acid in mol/l.
In a titration, 20 cm3 of potassium hydroxide (KOH) solution
was neutralised by 42.6cm3 of 0.5 mol/l hyrdochloric acid
(HCl).
Calculate the concentration of the base in mol/l.
In a titration, 2 mol/l sodium hydroxide (NaOH) solution was
neutralised by 22cm3 of 0.1mol/l sulphuric acid (H2SO4).
Calculate the volume of the base in cm3 that was neutralised.
In a titration, 0.05 mol/l potassium hydroxide (KOH) solution
was neutralised by 17.1cm3 of 0.25mol/l phosphic acid (H3PO4).
Calculate the volume of the base in cm3 that was neutralised.