Transcript The atomic structure - Chemistry Resources for IB, AP

The atom
Mass Spectrometer
Electron arrangement
Electron configuration
The atom
• State the position of protons, neutrons and electrons in the
atom.
• State the relative masses and relative charges of protons,
neutrons and electrons.
• Define the terms mass number (A), atomic number (Z) and
isotopes of an element.
• Deduce the symbol for an isotope given its mass number and
atomic number.
• Calculate the number of protons, neutrons and electrons in
atoms and ions from the mass number, atomic number and
charge.
• Compare the properties of the isotopes of an element.
• Discuss the uses of radioisotopes
State the position of protons, neutrons and electrons in the atom.
• Protons and neutrons are present in the
nucleus of an atom, electrons are in orbits
or shells around the nucleus.
• Atomic number, Z = number of protons; the
fundamental characteristic of an element.
Relative masses and relative charges of protons, neutrons and electrons.
• Relative masses: p = 1, n = 1, e = 1/1840;
charges: p = +1, n = 0, e- = -1.
The atom
• Isotopes are atoms of the same element
having the same atomic number / proton
number but different mass number/ number of
neutrons
• Isotopes differ in physical properties that
depend on mass such as density, rate of
diffusion etc.
• Chemical properties are the same because of
the same electronic configuration or
arrangement.
The atom
• Atomic mass of an atom is the average of the
atomic masses of its isotopes; depends on
isotopes relative abundance; leads to noninteger atomic masses.
• Atomic number is the number of protons in
the nucleus of an atom.
Radioactive isotopes of all elements can be
produced by exposing the natural element to a
flux of slow moving neutrons in a nuclear
reactor.
This results in the nucleus of the atom
capturing an additional
neutron.
Radiocarbon dating
•The rate of radioactive decay of carbon-14, can
be used to date objects.
•Naturally occurring carbon in living organisms
contains a fixed proportion of carbon-14 owing
to exchange with carbon in the atmosphere.
•On death this interchange stops and the
proportion of carbon-14 starts to decrease. After
about 5,700 years the proportion of carbon-14
will have fallen to about half its initial value.
• PEOPLE HAVE CLAIMED
THAT THIS SHROUD
WAS USED TO RAP THE
BODY OF JESUS.
• ISOTOPES STUDY GAVE
ANOTHER ANSWER
Radioisotopes as tracers
•Another use of radioisotopes is
as “tracers”. this relies on the fact
that the radioactive isotopes
behave chemically, and thus
biologically, in an identical
manner to the stable isotopes.
•The activity of the thyroid gland,
which preferentially absorbs
iodine, can be measured by
monitoring the increase in
radioactivity of the gland after
taking a drink containing traces of
iodine radioisotopes (typically 125I
and 131I).
Radioisotopes in radiotherapy
•Some radiosotopes
produce gamma rays and
hence can be a source of
quite intense radioactivity.
•`Cobalt-60 is an example of
this and radiation from
cobalt-60 sources is used in
radiation treatment for
cancer and industrially in
devices such as those
monitoring the thickness of
steel plate from a rolling mill.
Mass Spectrometer
• Describe and explain the operation of a mass
spectrometer.
• Describe how the mass spectrometer may be
used to determine relative atomic mass using
the 12C scale.
• Calculate non-integer relative atomic masses
and abundance of isotopes from given data.
Mass Spectrometer
Mass Spectrometer
• Stages of Operation:
–
–
–
–
–
Vaporization of sample
ionization to produce M+ ions
acceleration of ions by electric field
deflection of ions by magnetic field
detection of ions.
Mass Spectrometer
• Region A contains the vapourised
substance. If it is already a gas, then it
will contain the gas at low pressure, if
the sample is a solid or liquid, it must
be heated to produce the vapour.
Mass Spectrometer
• In region B, the particles are converted
from neutral atoms or molecules into
positive ions. This is usually done by
bombarding them with fast moving
electrons that are accelerated between
the two plates shown. These electrons
collide with electrons in the particle
knocking them out and leaving a
positive ion.
Mass Spectrometer
• In region C, these positive ions are
accelerated by the high electrical
potential difference between the two
parallel electrodes with holes in their
centers.
Mass Spectrometer
• In region D these fast moving ions
enter a magnetic field produced by an
electromagnet. The poles, shown as
circles, are above and below the plane
of the diagram. This causes the fast
moving ions to deflect
Mass Spectrometer
• Particles of a certain mass (dependent
on the field strength) will continue
round the tube and strike the detector
plate. Those with a greater mass will
not be deflected as much and those
with a smaller mass (deflection
depends on the charge to mass ratio
m/z).
Mass Spectrometer
• In region E particle will be detected by
means of the current flow required to
neutralise the positive charge that they
carry - the greater the number of
particles of a given mass that are
present, the greater the current.
Mass Spectrometer
• Degree of deflection:
• Lower the mass, higher the deflection.
• Higher the charge, higher the deflection.
• Deflection reflects mass/charge ratio; for
charge of +1, deflection depends on mass.
Mass Spectrometer
• For an element, the
mass spectrum
gives two important
pieces of
information: the
number of isotopes,
and the abundance
of each isotope;
thus the relative
average atomic
mass, Ar can be
calculated.
Mass Spectrometer
• For a molecule, the highest peak represents
the molecular (parent) ion and its mass gives
the relative molecular mass, Mr of the
compound (and the fragmentation pattern can
help determine its structure).
Electron Arrangement
• Describe the electromagnetic spectrum.
• Distinguish between a continuous spectrum
and a line spectrum.
• Explain how the lines in the emission
spectrum of hydrogen are related to electron
energy levels.
• Deduce the electron arrangement for toms
and ions up to Z = 20
Atomic emission spectrum
• The study of the emission of light by atoms
and ions is the most effective technique for
deducing the electronic structure of atoms.
• The term “light” is being used rather loosely
to indicate electromagnetic radiation. This
covers radiation from gamma rays through to
radio waves
Electromagnetic spectrum
• Electromagnetic waves have a range from
very low energy waves –like radio waves- and
very-high energetic waves like gamma
radiations.
Emission spectra
• Each element gives its characteristic set of
colours .
• These colours could be observed by a
spectrometer as lines each of fixed
wavelength called emission spectrum.
• Emission spectra are not continuous but
consist of separate lines.
• These lines become closer ( converge
towards the higher energy end of the
spectrum
Bohr atomic structure
• The electron travels in orbits around the nucleus of the atom.
• When an electron absorbs energy it moves to a higher level.
• When an electron moves down to a lower level it emits a packet of
energy called quantum.
• Each packet corresponds to a certain wavelength and shows a
certain colour.
• Continuous spectrum was not observed which meant that electrons
can only exist in specific levels but not in between them.
• The value of the energy level n=1,2,3….. Is called the principle
quantum number
Bohr Atomic model
• Electrons falling from the outer
levels to level 1 will emit the highest
amount of energy, and the spectra
will be in the UV region.
• Electrons falling to level 2 will form a
spectrum that falls in the visible
region.
• Electron falling to level 3 will form a
spectrum that falls in the infra red
region
n=3
n=2
n=1
Hydrogen visible spectrum
• The emission spectrum of hydrogen shows
four discrete lines in the visible region.
• The lines are converging towards the violet
which is more energetic.
• There are other line spectrum for hydrogen
in the invisible region
Convergence limit and ionization energy
•If sufficient energy is given to the atom,
it is possible to excite the electron
beyond the highest energy level.
•The electron will escape and the atom
will become an ion.
1. Why do these levels mean that the atom will show an
emission spectrum of discrete lines rather than a continuous spectrum.
2. Which three of the lettered energy changes involve absorption of energy by the atom?
3. Which three levels involves emission?
4. Of the three energy changes that involve emission, one results in blue light, one results in yellow light and the
third results in ultraviolet light
1. Which lettered change involve the emission of blue light?
2. Which lettered change involve the emission of yellow light?
3. Which lettered change involve the emission of ultraviolet light?
4
3
A
B
C
D
E
F
2
1
Emission Spectrum
• When electrons are excited, they jump to
higher energy levels.
• Electrons fall back to lower energy levels,
and the energy equivalent to the
difference in energy level is emitted in the
form of photons.
Emission Spectrum
• A continuous spectrum contains light of all
wavelengths in the visible range.
• A line spectrum consists of a few lines of
different wavelengths.
Emission Spectrum
• Energy levels come together in terms of
energy the farther away they are from the
nucleus; this explains the convergence
of lines in a line spectrum.
Electron configuration
• Explain how evidence from first ionization energies across
periods accounts for the existence of main energy levels and
sub-levels in atoms.
• Explain how successive ionization energy data is related to the
electron configuration of an atom.
• State the relative energies of s, p, d and f orbitals in a single
energy level.
• State the maximum number of orbitals in a given energy level.
• Draw the shape of an s orbital and the shapes of the px, py and
pz orbitals.
• Apply the Aufbau principle, Hund’s rule and the Pauli exclusion
principle to write electron configurations for atoms and ions up
to Z=54
Ionization energy
• The ionisation energy of an atom is the
minimum amount of energy required to
remove a mole of electrons from a mole of
gaseous atoms to form a mole of gaseous
ions.
• This change is endothermic because work is
needed to remove the electron.
Ionization energy
• Factors affecting the
magnitude of the
ionization energy:
–the charge on the nucleus.
–Shielding of electrons in Filled
inner orbitals.
–repulsion that the electron
experiences from other
electrons within the same
shell.
Ionization energy
Ionization energy
• Going down a group of the periodic table, the
ionisation energy of the elements decreases.
• This is because of a reduction in the amount
of electron-electron repulsion and hence the
greater nuclear-electron attraction that results
causes the remaining electrons to move
closer to the nucleus.
Ionization Energy
Ionization energy
• going across a period (for example period 2
from Li to Ne, or period 3 from Na to Ar), it can
be seen that the ionisation energy increases.
• This is because of the increase in the charge
on the nucleus which, as the electrons being
removed are all in the same energy level,
increases the effective nuclear charge, and
hence the ionisation energy.
Ionization energy
• The more electrons that have been removed
from an atom, the greater the energy required
to remove the next electron.
Electron configuration
• The maximum number of electrons in a main
energy level n is 2n2:
• 1st energy level, n = 1; maximum 2 e-;
• 2nd energy level n = 2, maximum 8 e-;
• 3rd energy level n = 3, maximum 18 e-.
Electron configuration
• The electron arrangement (or configuration)
indicates the number of electrons and their
energy distribution. This determines an
element’s physical and chemical properties.
Electron configuration
• Main energy levels, sub-levels and orbitals:
• The main energy levels, n are assigned
whole number integers, n = 1, 2, 3, 4… .
n = 1 represents the lowest energy level.
• Each main energy level contains n sublevels and a total of n2 orbitals.
• s, p, d, f etc. is the common notation for
sub-levels and orbitals within sub-levels.
Electron configuration
• An orbital is an area of space around the
nucleus in which an electron moves.
• Orbitals have characteristic shapes. There is
one s orbital which is spherical in shape,
three p orbitals which are dumbbell shaped,
called px, py pz, and arranged in the x, y, and
z directions respectively, five d orbitals and
seven f orbitals (both with complex shapes).
The relative energies of s, p, d, and f orbitals
with in a sub-level are: s < p < d < f.
Electron configuration
Electron configuration
• Each orbital can have a maximum of 2
electrons.
–n = 1 has one sub-level which is called an s sub-level
and which contains one s orbital.
–n = 2 has two sub-levels: 2s and 2p;
–n = 3 has 3 sub-levels: 3s, 3p and 3d;
–n = 4 has 4 sub-levels:4s, 4p, 4d and 4f, etc.
Electron configuration
• Each energy sub-level is
divided into orbitals
each of which can
contain up to two
electrons, which must
have opposite spins.
• The Pauli exclusion
principle, says that no
two electrons in an atom
can be in exactly the
same state (that is, they
cannot be in the same
place at the same time).
Electron configuration
• The electrons in atoms
always adopt the
lowest energy
configuration possible
by filling one sub-level
completely before
starting to fill the sublevel of next highest
energy. This is known
as the ‘Aufbau’
(building up) principle.
Electron configuration
• Hund’s rule, states
that sub-level
orbitals are singly
occupied as far as
possible by
electrons with the
same spin.
Electron configuration
• The orbitals are filled with electrons
according to the following order.
Electron configuration
o Which one of
the following
represents the
2p orbital of
Carbon
Electron configuration
• The arrangement of electrons in the porbital of some atoms is
Electron configuration
• The electronic structures of the
elements are related to the position of
the element in the periodic table.