Download Revision topic 1-3

Revision topic 1-3
Properties of gases:
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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
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Pressure is the amount of force exerted
on one unit of area.
Avogadro´s law
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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.
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This volume is known as the molar volume of a gas,
Vmolar (unit dm3/mol)
n = V / Vmolar
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/gasesv6.swf
Boyles law
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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
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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 collide
with the walls with more energy
and more frequency
P = k 2T
Charles' law
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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
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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.
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The ideal gas equation
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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
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Only ideal gases will follow
this equation exactly.
Isotopes
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Atoms of the same element always contain the same
number of protons, but if they contain a different
number of neutrons they are called isotopes.
• Isotopes have:
- The same number of protons = same atomic
number (Z)
- Different number of neutrons, thus different
mass number (A)
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Relative atomic mass
• The mass of an atom depends on the number of protons and
neutrons in the nucleus (the mass of electrons is so small that it
can be ignored in chemistry).
• The relative atomic mass is a weighted average mass
(according to relative abundances) of all the naturally occuring
isotopes of an element compared with an atom of the C-12
isoptope, which has a mass of exactly 12.
• e.g. Ar (H) = 1,008
• Many isotopes of elements are radioactive, because their nuclei break
down spontaneously and emit radiation.
2.2 The mass spectrometer
• The mass of individual isotopes can be
determined using a mass spectrometer:
1) sample is vaporized
2) ionized to positive ions: M (g) → M+ (g) + e3) the ions are accelerated and will pass through a
charged slit
4) the ions are deflected by an external magnetic
field
5) the ions are recorded on a detector (relative
amounts of the ions, mass to charge ratio)
• Electromagnetic radiation can be viewed as a stream
of tiny packets (quanta) of energy.
• The energy of electromagnetic radiation is expressed:
E = hf
where:
h = Planck's constant ≈ 6,626 · 10-34 J s
• The smaller the wavelength, thus the higher the
frequency, the more energy the wave possesses.
Line spectra
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When white light is passed through hydrogen gas, some of the
light is absorbed.
A continous spectrum is an emission spectrum that contains all
the wavelengths or frequencies of visible light.
An absorption spectrum is produced where some colours are
missing ( those that are absorbed by hydrogen).
A corresponding emission line spectrum has only certain
wavelegths or frequencies of visible light.The lines correspond
to the light of particular wavelengths given off (=emitted) by
the element.
Excitation
• When an atom absorbs energy,
an electron moves into an orbit
of higher energy further from
the nucleus. An unstable
excited state is produced.
• The electron soon falls back to
a lower level and gives out (=
emittes) the energy in the form
of electromagnetic radiation
with a specific frequency, f.
• The lowest energy level of the
electron is called the ground
state.
∆Eelectron = hf
Topic 3. Peridoicity
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Periodicity = repeating patterns of physical and chemical
properties.
3.2 Physical properties
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Groups: the group number (1-7) gives the number of
electrons in the outermost energy level (= valence electrons).
All atoms in the same group (or column) have similar
chemical properties.
Periods: All elements in the same period have the same
number of energy levels (electron shells).
Effective nuclear charge
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In every atom there is a balance between the attraction
of the positively charged nucleus for the negatively
charged electrons and repulsion between the electrons.
The outer electrons (valence electrons) do not
experience the full attraction of the positive nucleus
because of the presence of inner electrons. They are
shielded from the nucleus and repelled by the inner
electrons.
The effective charge experienced by the outer
electrons is less than the full nuclear charge.
Atomic radius
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The effective charge increases (=
atomic radii decreases) as a period
is crossed from the left to the right:
● One proton is added to the nucleus
and one electron to the outermost
electron shell.
● There is no change in the number
of inner electrons.
The effective charge remains almost
the same down a group, because:
● The increase in the nuclear charge
is offset by the increase in the
number of inner electrons.
Ionic radius
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Positive ions are smaller than their parent atoms (because
of loss of the outer shell).
Negative ions are larger than their parent atoms (because
of increased electron repulsion by addition of electrons).
The ionic radii decrease as a period is crossed from the
left to the right (because of increased attraction between
the nucleus and the electrons).
The ionic radii increase down a group as the number of electron
shells increase.
Electronegativity
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A measure of the tendency of an atom in a molecule to
attrat a pair of shared electrons towards itself.
Electronegativity values increase from left to right
across the periodic table and decrease down a group.
The first ionization energy
• The energy required to remove one mole of electrons from one
mole of isolated gaseous atoms of an element to form one mole of
gaseous unipositive ions under standard conditions.
X (g) → X+ (g) + e●
The second and third ionization energies : X+ (g) → X2+ (g) + eX2+ (g) → X3+ (g) + e-
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The stronger the bonds, the higher the melting point.
3.3 Chemical properties
Group 1: the alkali metals
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Shiny, silvery, soft
Reactivity increases down the group, because the first
ionization energy decrease down a group.
https://www.youtube.com/watch?v=uixxJtJPVXk
Group 7: the halogens
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Reactivity decreases down a group as the atomic radius increases
and the attraction for outer electrons decreases.
Halide formation: 2 Na (s) + Cl2 (g) → 2NaCl (s)
Displacement reactions: a more reactive halogen can displace the
ions of a less reactive halogen (= take the electron from the less
reactive ions) .
Ex. Which of the following chemical reactions are possible?
d
Chemical properties of elements in the
Period 3
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Na, Mg and Al are metals. They are shiny and good
conductors of heat and electricity.
Si is a semi-conductor and is called a metalloid since it
has some of the properties of a metal and some of a
non-metal.
P, S, Cl and Ar are all non-metals and do not conduct
electricity.
The Period 3 oxides
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Metals can also be distinguished from non-metals by their chemical
properties.
Oxides are products of reactions in which an element reacts with
oxygen.
Metal oxides
Metal oxides (Na2O, MgO) are ionic compounds:
- solids in room temperature
- high mp & bp
- conduct electricity when molten (or in aqueous
solutions)
- basic in aqueous solutions:
Na2O (s) + H2O (l) → 2 NaOH (aq)
MgO (s) + H2O (l) → Mg(OH)2 (s)
Aluminium oxide
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Al2O3 is an ionic oxide with some covalent character.
It is amphoteric as it acts as a base when it reacts with
acids and acts as an acid when it reacts with bases:
Silicon oxide
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Silicon oxide, SiO2, has a giant covalent structure with
very high melting and boiling points.
It is insoluble in water.
It is classified as an an acidic oxide, because it reacts
with NaOH at temperatures above 350° C.
Non-metallic oxides
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Nonmetal oxides are covalently bonded because of the
small difference in the elements' electronegativity values.
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sulfur: SO2, SO3
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chlorine: Cl2O, Cl2O7
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phosphorus: P4O6, P4O10
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They have low mp and bp
and do not conduct electricity.
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Non-metallic oxides are acidic in aqueous solutions:
P4O10 + 6 H2O (l) → 4 H3PO4 (aq)
phosphoric(V)acid
SO3 (g) + H2O (l) → H2SO4 (aq)
sulfuric(VI)acid
Acidic rain
• Rain is naturally acidic, because the water molecules react
with the CO2 in the air and form the weak acid H2CO3.
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Acid rain is therefore: precipitation (rain, snow) with pH
lower than 5.6.
The main acids present in acid rain are sulfuric acid nitric
acid.
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The main acids present in acid rain are sulfuric acid nitric acid.
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The sulfuric acid in the rain reacts with calcium carbonate (in
limestone or marble) to create calcium sulfate, which then flakes
off.
CaCO3(s) + H2SO4(aq) → CaSO4(aq) + CO2(g) + H2O(l)
TOPIC 2 (HL)
● Graph of successive ionization energies for sodium :
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Sub-levels of electrons (HL)
• 1. Aufbau Principle: electrons are placed into orbitals of lowest
energy first.
• 2. Pauli exclusion principle: no two electrons in an atom can
have exactly the same four quantum numbers
no more
than two electrons can occupy any one orbital, and if two
electrons are in the same orbital they must have the opposite
spin.
1s
• 3. Hund´s rule: If more than one orbital in a sub-level is
available, for example px, py, pz , electrons occupy the orbitals
single with parallel spin before they are paired up.
/watch?v=sMt5Dcex0kg
The shapes of s, p and d orbitals
The shapes of s, p and d orbitals
Halv-filled and filled sub-levels are more stable:
Cr
Cu
Electron configuration
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Electron arrangement: the number of electrons in each
main level (e.g. Na: 2, 8,1)
Electron configuration: the number of electrons in each
sub-level (e.g. Na: 1s22s22p63s1)
When writing electronic configurations, the sum of the
superscripts must always total the number of electrons
in the atom (or ion).
13.2 First row d-block elements (HL)
Physical properties
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Small atomic radii compared to the neighbouring sblock elements.
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Only a small increase in atomic radii across the period.
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Metallic bonding
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High density
Chemical properties
Transition metals
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An element that contain an incomplete d sub-level in
one or more of its oxidation states
= forms ions with partially filled d sub-levels
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Neither Zn nor Sc are considered to be transition
elements:
Transition metals
Complexes
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Because of their small size and relatively high charge, the
transition metal ions have a high charge density.
They attract species that are rich in electrons: ligands.
A ligand is a molecule or negative ion that donates a pair
of electrons to a central metal ion to form a dative (or
coordinat) covalent bond, e.g. H2O, NH3, Cl-, CN-
[Fe(H2O)6]3+
Ligand exchange (replacement)
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In aqueous solutions the water molecule usually
act as a ligand, but it can be replaced with
another ligand.
Transition metals absorb visible light
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In an isolated gaseous transition metal atom, the five 3 d
orbitals are degenerate since they all have the same energy.
However, since the 3 d sub-shells all have different orientations
in space they will be orientated differently relative to the
ligands in a complex ion.
The 3d electrons close to a ligand will experience repulsion and
be raised in energy.
The 3d electrons further from the ligand will be reduced in
energy.
→ The 3d sub/shell splits into two energy levels.
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If all light is absorbed, the substance appears black.
If only certain wavelengths are absorbed, the compound
appears coloured.
If all light is reflected, the compound appears white.
Heterogeneous catalysts
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The catalyst is in a different state from the reactants.
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e.g. Fe in the Haber process (ammonia is produced)
Palladium and platinum in converters in cars
Homogenous catalysts
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Homogenous catalysts are in the same phase as the
reactants and products.
The two reacting species bond chemically to the
transition metal for ing an intermediate, and then leave.
Transition metals can be relatively easily oxidized and
reduced due to their variable oxidation states.