Download IB 1 CHEMISTRY

IB 1 CHEMISTRY
Teacher: Annika Nyberg
[email protected]
Course book:
Pearson Baccalaureate:
Higher Level Chemistry for the IB Diploma
2nd edition ISBN10: 1447959752
isbn13: 9781447959755
Pearson Baccalaureate:
Standard Level Chemistry for the IB Diploma
2nd edition
ISBN 10: 1447959760
Isbn13: 9781447959762
http://www.chem1.com/acad/webtext/virtualtextbook.html
1. Stoichiometric Relationships
NOS: The atomic theory
●
●
●
All matter is composed of atoms.
These atoms cannot be created or destroyed during
chemical reactions, they can only be rearranged.
Physical and chemical properties of matter depend on
the bonding and arrangement of these atoms.
NOS: The phlogiston theory
https://www.youtube.com/watch?v=3MuMPLoQZN4
The atom
• Atoms are composed of subatomic
particles: protons, neutrons and electrons.
• Protons and neutrons are found in the nucleus of
the atom and are therefore called nucleons.
In a neutral atom, the number of electrons
equals the number of protons.
1.1 Matter
●
The characteristics of matter:
has a mass
➢ occupies a volume in space
➢ is made up of particles (atoms, moleules or
ions)
➢ particles are in constant motion
➢
States of matter
• Substances have different appearances and physical
properties in different states of matter.
http://www.youtube.com/watch?v=s-KvoVzukHo&feature=related
Temperature
●
The way the particles
of matter move
depends on the
temperature.
temperature (K) = temperature (ºC) + 273,15
Changes of state
http://www.youtube.com/watch?v=6s0b_keOiOU
Ice melting
http://www.youtube.com/watch?v=pP_lZaOchE0
Dry ice bubbles
Heating curve for water
The philosophers stone
https://www.youtube.com/watch?v=zTsoClR49UI
Greatest discoveries chemistry
https://www.youtube.com/watch?v=s7xxMX4Ovig
Elements and compounds
●
●
●
An element contains atoms of only one type.
Compounds are made up of more than one
element, chemically bonded together.
The properties of a compound is very different
from those of its constituent elements, e.g. NaCl.
+
→
Mixtures
●
●
Elements and compounds have a constant
composition and are pure substances.
Pure substances can combine to form a mixture.
• Mixtures can be either:
The language of chemistry
●
●
●
Chemistry has a universal language.
The International Union of Pure and Applied Chemistry
(IUPAC) is an organization that develops a system of
standardized nomenclature for chemical compounds.
See the IUPAC gold book for chemical terminology:
http://goldbook.iupac.org/index.html
氯化钠
‫سدیم کلرید‬
Cloruro di sodio
http://en.wikipedia.org/wiki/Sodium_chloride
http://fa.wikipedia.org/wiki/%D8%B3%D8%AF%DB%8C%D9%85_%DA%A9%D9%84%D8%B1%DB%8C%D8%AF
Writing equations
●
●
●
●
Write the correct formulas for all the reacting species,
reactants on the left-hand side and the products on the righthand side.
Write the correct coefficients in front of each species. The
reaction is then said to be stoichometrically balanced.
It is good practice to include the state symbols: (s), (l), (g),
(aq).
Ions that remain unreacted in the reaction (= spectator ions)
can be left out from the equation.
Ionic compounds
●
In forming ionic compounds the number of ions
used is such that the number of positive charges
is equal to the number of negative charges → the
compound is electrically neutral, e.g.
copper(II)carbonate
sodium sulfate
Molecules
●
Some elements exist as molecules = one or
more atoms held together by covalent bonds.
The atom economy
●
In an ideal chemical process the amount of
products equals the amount of reactants and no
atoms are “wasted”.
percentage
=
atom economy
Moleculas mass of useful products
Moleculas mass of atoms in reactants
https://www.youtube.com/watch?v=Zuyk4hfbjSA
x 100%
1.2 The mole concept
●
●
SI: the international system of measurment.
The SI (Systeme International d´Unités) system
has seven base units. All the other units are
derived from them.
The gram was originally defined in 1795
as the mass of one cubic centimeter of
water at 4 °C, making the kilogram equal
to the mass of one liter of water. The
prototype kilogram, manufactured in
1799 and from which the current
kilogram is based, has a mass equal to the
mass of 1.000025 liters of water.
●
●
A single atom of an element has an extremely small
mass, far too small to weigh.
Ex. Calculate the mass of one hydrogen-1 atom.
1,67 ▪ 10-27 kg (+ 9,1 10-31 kg) =
0,00000000000000000000000167 g
●
Chemists measure amount of substances in moles, by
counting particles (atoms, molecules or ions).
Amount of a substance, n
●
●
12,00 grams of carbon-12 contains 6,02 ·1023
atoms. This number is called the Avogadro's
constant (NA or L).
One mole of ANY substance contains 6,02 ·1023
particles.
Isotopes
●
●
•
•
•
The atoms of an element are made up of an
mixture of isotopes.
Isotopes have:
- the same atomic number (= the same number of
protons and electrons)
- different mass number (different number of
neutrons)
- similar chemical properties, but different physical
properties (density, bp)
Relative atomic mass, Ar (no unit)
●
●
●
The relative atomic mass of an element is the
weighted mean mass of all the naturally occuring
isotopes of that element relative to the mass of
carbon-12.
For example, hydrogen has 1/12 of the mass of
carbon-12.
1 mol of hydrogen atoms has a mass of 1.01 g.
Sub-atomic particle
Mass (kg)
Relative mass
proton
1,67 ▪ 10-27 kg
1
neutron
1,67 ▪ 10-27 kg
1
electron
9,1 ▪ 10-31 kg
0,0005
Relative molecular mass, Mr
●
●
The relative molecular mass, Mr, of a molecule
is the sum of the relative atomic masses of the
atoms in the molecule (found in the periodic
table).
Mr (C2H5OH) =
The relative formula mass, Mr, is similar to the
above but can be used with nonmolecular
substances such as ionic compounds.
Mr (AgNO3) =
Molar masses, M (g mol-1)
●
●
The mass of one mole of ANY substance is
known as the molar mass.
For example, 1 mol of iron contains 6.02 x 10 23
iron atoms and has a mass of 55.85 g.
602 000 000 000 000 000 000 000
Formulas of compounds
●
●
●
The empirical formula is obtained experimentally by burning
a compound in oxygen so that all its elements forms oxides.
The amount of oxides can be determined and that gives the
original amount of each element.
The empirical formula shows the simplest whole number
ratio of atoms of each element in a particle of that substance,
e.g. C6H12O6
Formulas of compounds
●
●
●
The empirical formula is obtained experimentally by
burning a compound in oxygen so that all its
elements forms oxides.
The amount of oxides can be determined and that
gives the original amount of each element.
The empirical formula shows the simplest whole
number ratio of atoms of each element in a particle
of that substance, e.g. CH2O
Molecular formula:
Structural formula:
●
Remember to check that all the percentages add
up to 100%!!
Ex. 17
Calculate the empirical formula of a compound that contains
40.4% carbon, 6.0 % hydrogen, 17.7% nitrogen and 35.9%
oxygen by mass.
●
Ex.18
A compound contains 12.79% carbon, 2.15%
hydrogen and 85.06% bromine by mass. Its
relative molecular mass is 187,9. Determine a) the
empirical formula and b) the molecular formula of
the compound.
g)
h)
i)
j)
1.3 Reacting masses and volumes
●
●
Chemical equations show reactants combining in
fixed ratios (moles) to form products.
Ex. Methane burns in air:
CH4 (g) + 2 O2(g) → CO2 (g) + 2H2O (g)
●
Calculate the mass of carbon dioxide produced
from the complete combustion of 1,00 g of
methane.
Theoretical yield
●
The balanced chemical equation can be used to
predict how much product can theoretically be
produced from given masses of starting material.
Ex 7.
a) Write the balanced chemical equation for the
reaction where ethene and steam react to produce
ethanol (C2H5OH).
b) What is the maximum amount of ethanol that can
be produced when 1,0 kg of ethene and 0,010 kg of
steam are placed into the reaction vessel?
Percentage yield
1.5 Solutions
●
●
Some liquids are pure substances, but more
commonly liquids are solutions containing two
or more components.
Solution: homogenous mixture of two or more
substances.
solid/solid:
● solid/liquid:
● liquid/liquid:
● gas/liquid:
●
Concentration
●
The concentration of a solution is the amount of
solute (in moles) per volume of solution.
Dilution of solutions
●
Stock solution: a concentrated starting solution
Density
density = _mass
volume
Acid-base titration
Ex. 12 Acid-base titrations:
●
Sodium hydroxide reacts with hydrochloric acid
according to the following equation:
NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l)
●
Calculate the volume of 0,0500 mol dm-3 sodium
hydroxide solution to react exactly with 25 cm3 of
0,20 mol dm-3 hydrochloric acid.
1.4 Mass and gaseous volume
relationships in chemical
reactions
Properties of gases:
●
●
●
●
●
Gases have a small mass
All gases respond in a similar way to
changes in temperature, pressure
and volume.
They exert a pressure, that depends
on the amount of gas and the temperature
There is no bonding between molecules
The molecules may move in all directions allowing
the gas to expand throughout any container
Pressure
●
Pressure is the amount of force
exerted on one unit of area.
Avogadro´s law
●
One mole of any gas will occupy the
same volume, if the temperature
and pressure are the same.
= equal volumes of different gases at the
same temperature and pressure contain
the same number of particles.
n = V / Vmolar
●
This volume is known as the molar volume
3
●
●
At standard temperature and pressure, STP (273
K and 101,3 kPa) one mole gas occupies 22,4
dm3.
At room temperature, RTP (298 K and 101,3
kPa) one mole gas occupies 24 dm3.
●
●
Ex. 16 s. 20: Calculate the volume
occupied by 4,40 g of carbon dioxide at
STP.
Worked example, s.21: What volume of
hydrogen is produced when 0,056 g of
lithium reacts completely with water at
STP:
2Li (s) + 2 H2O (l) 2LiOH (aq) + H2 (g)
●
Ex. 17 s. 21: Calcium reacts with water to
Ex 6.
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/gasesv6.swf
Boyles law
●
●
PV = k1
i.e. the relationship between volume and
pressure for a gas
The law describes how the volume of a given
amount of gas at constant temperature varies
inversely with the applied pressure:
Gay Lussac's law
●
●
i.e. the relationship
between temperature
and pressure for a gas
states that the pressure
of a given amount of gas
held at constant volume
is directly proportional to
its temperature in kelvin
= an increase in
temperature increases
the kinetic energy of the
particles, which means
they will move faster and
P = k 2T
Charles' law
●
●
i.e. the effect of
temperature on the
gas volume
The law describes how
the volume of a given
amount of gas is
directly proportional to
its temperature in
kelvins.
V = k3T
The combined gas law
●
The three gas laws can be combined to one expression:
where 1 refers to the initial conditions and 2 to the final
conditions.
Only ideal gases obey the gas laws perfectly. Real
gases can be treated as ideal gases, unless dealing
with extremely precise measurements.
●
Ex 7.
The behaviour of gases
●
●
There is large intermolecular distance
between the molecules in a gas = a gas
is mainly empty space.
The gas molecules are free to move
randomly in all directions. The
molecules travel in a straight line until
they collide with other gas molecules
or with the walls of the container.
Kinetic theory of gases
●
●
●
There are no attractive or repulsive forces
between the atoms or molecules in a gas.
The kinetic energy of a gaseous molecule or
atom is given by the expression E = ½ mv2,
where m is the mass of the particle and v is its
speed.
The mean kinetic energy of molecules or atoms
in a gas is directly proportional to its temperature.
Ideal gas
●
●
●
A gas that obeys all the gas laws under all conditions
is said to be an ideal gas.
No real gas behaves exactly as an ideal gas, but at
high temperatures and low pressures the model
describes most gases well.
●
At high pressure and low temperature the gas
particles are compressed:
- the gas particles move slowly and come so close
together that intermolecular forces attract them to
each other
- the actual volume occupied by the particles will be
significant compared to the total volume of the gas.
The ideal gas equation
●
The ideal gas law relates pressure, volume,
temperature and amount of substance:
where R is the gas constant
= 8,31451 J K-1 mol-1
●
Only ideal gases will follow
this equation exactly.
Ex. 8. At 273 K and 101 325 Pa, 12,64 grams of
3
a gas occupy 4,00 dm . Calculate the molar mass
and relative molecular mass of the gas.
Ex 9.
Ex 10.
Ex 11.
●
●
Aim 8: The negative environmental impacts of
refridgeration and air conditioning systems are
significant.
The use of CFCs as refridgerants has been a
major contributor to ozone depeltion.
https://www.youtube.com/watch?v=XLY8m-dXOxo
https://www.youtube.com/watch?v=c5mfbm9oREk&list=PLgGh_MSicdaHo8
OwhRxyC5I3YvbWAYyKL&index=5