Transcript Chapter 13

Electrons in Atoms
Democritus (400 B.C.)
• Proposed that matter was
composed of tiny indivisible
particles
• Not based on experimental
data
• Greek: atomos
Alchemy (next 2000 years)
• Mixture of science and mysticism.
• Lab procedures were developed, but alchemists
did not perform controlled experiments like true
scientists.
John Dalton (1807)
 British Schoolteacher
 based his theory on others’
experimental data
 Billiard Ball Model
 atom is a
uniform,
solid sphere
Henri Becquerel (1896)
 Discovered radioactivity
 spontaneous emission of
radiation from the nucleus
 Three types:
 alpha () - positive
 beta () - negative
 gamma () - neutral
J. J. Thomson (1903)
 Cathode Ray Tube
Experiments
 beam of negative particles
 Discovered Electrons
 negative particles within the
atom
 Plum-pudding Model
J. J. Thomson (1903)
Plum-pudding Model
 positive sphere
(pudding) with
negative electrons
(plums) dispersed
throughout
Ernest Rutherford (1911)
 Gold Foil Experiment
 Discovered the nucleus
 dense, positive charge in the
center of the atom
 Nuclear Model
Rutherford’s Gold Foil Experiment
(a) The results that the metal foil experiment would
have yielded if the plum pudding model had been
correct. (b) Actual results.
Ernest Rutherford (1911)
 Nuclear Model
 dense, positive nucleus surrounded
by negative electrons
Niels Bohr (1913)
 Bright-Line Spectrum
 tried to explain presence of
specific colors in hydrogen’s
spectrum
 Energy Levels
 electrons can only exist in
specific energy states
 Planetary Model
Niels Bohr (1913)
Bright-line spectrum
 Planetary Model
 electrons move in
circular orbits within
specific energy levels
Erwin Schrödinger (1926)
 Quantum mechanics
 electrons can only exist in
specified energy states
 Electron cloud model
 orbital: region around the
nucleus where e- are likely
to be found
Erwin Schrödinger (1926)
Electron Cloud Model (orbital)
 dots represent probability of finding an enot actual electrons
James Chadwick (1932)
 Discovered neutrons
 neutral particles in the
nucleus of an atom
 Joliot-Curie Experiments
 based his theory on their
experimental evidence
James Chadwick (1932)
Neutron Model
 revision of Rutherford’s Nuclear Model
Electromagnetic Radiation
 Electromagnetic radiation – radiowaves, X-rays,
microwaves, infrared waves, visible light,
ultraviolet waves and gamma rays.
 All electromagnetic radiation travel at the
speed of light (c = 3.0 x 108 m/s) in a vacuum.
The different wavelengths of
electromagnetic radiation.
The electromagnetic spectrum.
Physics and the Quantum Mechanical
Model
 Amplitude – wave’s height from the origin to
the crest.
 Wavelength (l)– distance between the crests.
 Frequency (u)– number of wave cycles to pass a
given point per unit of time.
Water wave (ripple).
Physics and the Quantum Mechanical
Model
 Frequency and wavelength are inversely
proportional. As frequency increases,
wavelength decreases, and vice versa, but their
product will always equal the speed of light.
c = lu
 SI units for frequency are cycles per second is a
hertz (Hz), or 1/seconds (1/s or s-1).
Relationship Between Wavelength
and Frequency
Physics and the Quantum Mechanical
Model
 What is the frequency of light that has a wavelength
of 550 nm? (1m = 109 nm or 1 nm = 10-9 m)?
 What is the wavelength of light, in cm, that has a
frequency of 9.60 x 1014 Hz (1/s)?
 What is the frequency of light (Hz) that has a
wavelength of 740 nm (1m = 109 nm or
1 nm = 10-9 m)?
Physics and the Quantum Mechanical
Model
 Sunlight splits into a spectrum of colors when it
passes through a prism.
 Colors of the spectrum include red, orange, yellow,
green, blue, indigo and violet.
 Red light has the longest wavelength and the
lowest frequency, while violet light has the shortest
wavelength and the highest frequency.
Dispersion of White Light By a
Prism
A photon of red light (relatively long wavelength)
carries less energy than a photon of blue light
(relatively short wavelength) does.
Physics and the Quantum Mechanical
Model
 Every element emits light after it absorbs energy.
The light that is emitted (atomic emission spectra)
is different for every element, and differs from
white light because it is not continuous.
 Max Planck said that color changes can be
explained if you assume that the energy of a
substance changes in small increments.
Emission (line) Spectra of Some
Elements
Emission (line) Spectra of Some
Elements (cont’d)
Emmision (line) Spectra of
Some Elements (cont’d)
Physics and the Quantum Mechanical
Model
 Planck showed that the amount of radiant energy
(E) absorbed or emitted by a substance is
proportional to the frequency of the radiation.
 E = hu
 h is Planck’s constant (6.626 x 10-34 J s)
 Any attempt to increase or decrease the energy of a
system by a fraction of h times u will fail because
energy is only emitted or absorbed in quanta, or
bunches of energy.
Planck’s Constant Examples
 What is the energy of a photon with a frequency of
2.94 x 1015 cycles per second (s-1 or Hz)?
 What is the energy of a light particle with a
wavelength of 675 nm?
Homework Problem Examples
 What is the wavelength, in nm, of light with a frequency
of 9.5 x 109 s-1? ( 1 m = 109 nm)
 How much energy is contained in a photon with a
wavelength of 5.17 x 10-4 m?
Planck’s Revelation
 Showed that light energy could be thought of
as particles for certain applications
 Stated that light came in particles called
quanta or photons
 Particles of light have fixed amounts of energy
 The energy of the photon is directly
proportional to the frequency of light
 Higher frequency = More energy in photons
Physics and the Quantum Mechanical
Model
 Photons – light energy. The energy of photons
is quantized according to the equation E = hu.
 Light was therefore thought to have a dual
wave-particle behavior to explain all of its
characteristics.
Electromagnetic radiation (a beam of light) can be
pictured in two ways: as a wave and as a stream of
individual packets of energy called photons.
Bohr’s Model
 Energy of an electron is related to the
distance electron is from the nucleus
 Energy of the atom is quantized
 atom can only have certain specific energy
states called quantum levels or energy
levels
 when atom gains energy, electron “moves” to
a higher quantum level
 when atom loses energy, electron “moves” to
a lower energy level
 lines in spectrum correspond to the
difference in energy between levels
Bohr’s Model
 Atoms have a minimum energy called the ground
state
 The ground state of hydrogen corresponds to having
its one electron in an energy level that is closest to the
nucleus
 Energy levels higher than the ground state are called
excited states
 the farther the energy level is from the nucleus, the higher its
energy
 To put an electron in an excited state requires the
addition of energy to the atom; bringing the electron
back to the ground state releases energy in the form of
light
(a) A sample of H atoms receives energy from
an external source.
(b) The excited atoms (H) can release the
excess energy by emitting photons.
When an excited H atom returns to a lower energy
level, it emits a photon that contains the energy
released by the atom.
Hydrogen atoms have several excitedstate energy levels.
Each photon
emitted by an
excited
hydrogen atom
corresponds to a
particular
energy change
in the hydrogen
atom.
Bohr’s Model
 Distances between energy levels decreases as
the energy increases
 light given off in a transition from the second
energy level to the first has a higher energy
than light given off in a transition from the
third to the second, etc.
 1st energy level can hold 2 electrons (e-1), the
2nd 8e-1, the 3rd 18e-1, etc.
 farther from nucleus = more space = less
repulsion
Models of the Atom
 Energy level – region around the nucleus where
the electron is likely to be found. Think of
steps on a ladder.
 Essentially, you must be in one energy level or
another, you can’t be between energy levels,
just like you can’t stand in mid-air between the
steps of a ladder.
The difference between continuous and quantized
energy levels can be illustrated by comparing a flight of
stairs with a ramp.
Models of the Atom
 Energy levels are not equally spaced. The
further away an electron is from the nucleus,
the easier it becomes to pull that electron off of
that particular atom.
 Erwin Schrodinger – in 1926, he came up with a
new way of describing the energy and location
of an electron, called the quantum mechanical
model, which is a mathematical method.
Models of the Atom
 The quantum mechanical model does not say
that electrons take exact paths around the
nucleus, but that it estimates the probability
(likelihood) of finding an electron in a certain
position.
 If the electron cloud is very dense, it is more
likely that you will find the electron there, then
if the electron cloud is less dense.
The probability
map, or orbital,
that describes
the hydrogen
electron in its
lowest possible
energy state.
Orbitals
 Orbital – area where an electron is likely to
be found.
 usually use 90% probability to set the limit
 three-dimensional
 Orbitals are defined by three integer terms
called the quantum numbers.
 Each electron also has a fourth quantum
number to represent the direction of spin
Models of the Atom
 Principal quantum number (n) – designates
the energy level of the electrons. n will always
be an integer.
 The distance from the nucleus increases as n
increases.
 Within each energy level, electrons occupy
energy sublevels.
 The number of energy levels (n) is always the
same as the number of sublevels.
Models of the Atom
 Sublevel – part of an energy level.
 1st energy level has 1 sublevel (“s” sublevel)
 2nd energy level has 2 sublevels (“s” and “p”
sublevels)
 3rd energy level has 3 sublevels (“s”, “p”, and
“d” sublevels)
 4th energy level has 4 sublevels (“s”, “p”, “d”
and “f” sublevels)
An illustration of how principal levels can be
divided into sublevels.
Principal level 2
shown divided
into the 2s & 2p
sublevels.
Models of the Atom
 Atomic orbitals – areas where electrons are
likely to be found.
 s orbital – spherical in shape, only 1 s orbital
per sublevel.
 p orbital – dumbbell shaped, 3 p orbitals per
sublevel.
 d orbital – 5 d orbitals per sublevel.
 f orbital – 7 f orbitals per sublevel.
The relative sizes of the
1s and 2s orbitals of hydrogen.
The 2p orbitals.
The five 3d orbitals.
Models of the Atom
 In any orbital, there can be a maximum of two
electrons.
 The maximum number of electrons that can
occupy an energy level is given by the formula
2n2, where n is the # of the energy level.
 1st energy level up to 2 electrons
 2nd energy level up to 8 electrons
 3rd energy level up to 18 electrons
 4th energy level up to 32 electrons
Quick Review
 Max of 2 electrons per orbital
 “s” sublevel – 1 orbital per sublevel (up to 2 total
electrons)
 “p” sublevel – 3 orbitals per sublevel (up to 6
total electrons
 “d” sublevel – 5 orbitals per sublevel (up to 10
total electrons)
 “f” sublevel – 7 orbitals per sublevel (up to 14
total electrons)
The orbitals being filled for elements in
various parts of the periodic table.
Electron Arrangements in Atoms
 Electron configuration – the way in which
electrons are arranged in energy levels outside
of the nucleus.
 Orbital notation – a way of showing the
electron configuration using arrows to
represent each electrons and boxes to represent
each orbital.
Electron Arrangements in Atoms Rules
 Aufbau principle – electrons enter orbitals of
lowest energy first.
 Pauli exclusion principle – an atomic orbital
may hold at most two electrons.
 Electrons within the same orbital have opposite
spins.
 Hund’s rule – one electron must be put in each
orbital of a sublevel before any one orbital can
have two electrons in it.
Orbital Notations
 When writing orbital notations, use one arrow to
represent each electron.
 Electrons must enter the lowest energy sublevel
possible before moving to a higher energy sublevel
 Even if you don’t have enough electrons to fill each
orbital of a sublevel, you must still show that those
orbitals exist.
 The total number of arrows (electrons) must be
equal to the atomic # for each element.
Types of Electrons in Arrangements
 Shared electrons – orbitals where there are two
electrons (arrows) with opposite spins.
 Unshared electrons – when an orbital only has
one electron in it.
 Shared pair of electrons – any orbital that
contains two electrons.
A box diagram
showing the
order in which
orbitals fill to
produce the
atoms in the
periodic table.
Each box can
hold two
electrons.
Orbital Notations
 Write the orbital notation for oxygen.
 Write the orbital notation for aluminum.
 Write out the orbital notation for cobalt
The orbitals being filled for elements in
various parts of the periodic table.
Electron Arrangements in Atoms
 Electron Configurations – the way in which
electrons are arranged around the nucleus of
an atom. Each configuration has 3 parts:
2
1s
 “1” represents the energy level, “s” represents
the sublevel, and “2” represents the number of
electrons in that sublevel
 The total of superscripts is equal to the atomic
number for the element.
Electron Arrangements in Atoms
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
The orbitals being filled for elements in
various parts of the periodic table.
Electron Configurations
 Which element is represented by the following
electron configuration:
 1s22s22p63s23p6
 1s22s22p63s23p64s23d104p65s24d105p66s2
4f145d106p67s1
Electron Configurations
 Write the electron configuration for the
following elements:
 Sulfur
 Gallium
 Thorium
 Platinum
Electron Configurations
 What is wrong with each of the following electron
configurations?
 1s22s22p63s23p63d104s24p5
 1s22s22p63s23p64s23d104p65s24d105p66s25d106p3
 1s22s22p63s23p64s23d84p65s1
Noble Gas Configurations
 Noble gas configurations are used as a shorthand
for long electron configurations.
 Find the noble gas before the element you are
writing the configuration for, put it in brackets,
and then start with the next s sublevel to fill out
the rest of the configuration.
The orbitals being filled for elements in
various parts of the periodic table.
The periodic table with atomic symbols,
atomic numbers, and partial electron configurations.
Noble Gas Configurations
 Write the noble gas configuration for the
following elements:
 Sulfur
 Iron
 Thorium
 Platinum
Noble Gas Configurations
 What element is represented by the following
noble gas configuration:
 [Kr]5s24d105p2
 [Ar]4s2
 [Xe]6s24f145d6
Noble Gas Configurations
 What is incorrect about the following noble
gas configurations?
 [Ar]2s22p2
 [Kr]4d10
 [At] 7s24f146d7
Electron Configuration
 Elements in the same column on the Periodic Table
have
 Similar chemical and physical properties
 Similar valence shell electron configurations
 Same numbers of valence electrons
 Same orbital types
 Different energy levels
 Valence electrons – outer energy level “s” and “p”
sublevel electrons or electrons that are furthest away
from the nucleus
Noble Gas Configurations & their
relation to the Periodic Table
 Lithium – [He]2s1
 Fluorine – [He]2s22p5
 Sodium – [Ne]3s1
 Chlorine – [Ne]3s23p5
 Potassium – [Ar]4s1
 Bromine-[Ar]4s23d104p5
 Rubidium – [Kr]5s1
 Iodine-[Kr]5s24d105p5
s1
1
2
3
4
5
6
7
s2
p1 p2 p3 p4 p5 s2
p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
Periodic Trends
 Atomic radius – distance from the nucleus of
an atom to its valence electrons. The radius
tells that size of the atom.
 Moving from left to right across a period,
atomic radius decreases.
 Electrons within the same energy level don’t
have as great of an effect on one another as
electrons from different energy levels.
Trend in Atomic Size
 Increases down column
 valence shell farther from nucleus
 Decreases across period
left to right
 adding electrons to same valence shell
 valence shell held closer because more
protons in nucleus
Periodic Trends
 Moving down a group, atomic radius increases.
 The valence electrons get further and further
from the nucleus because you are adding more
energy levels. Therefore the radius of the atom
increases.
Representation of Atomic Radii of the
Main-Group Elements
Periodic Trends
 Example: Put the following elements in
order of increasing atomic radius:
Zn, Sc, Se, K, Cs, O
 Example: Put the following elements in
order of decreasing atomic radius:
F, Cd, Ba, Ge, W, Cl
The orbitals being filled for elements in
various parts of the periodic table.
Periodic Trends
 Ionization energy – the energy required to
remove an electron from an atom (1st
ionization energy).
 Removing an electron creates a charge
imbalance, so a cation (positive ion) is
formed.
 2nd Ionization energy – the energy required
to remove two electrons from an atom.
Periodic Trends
 Moving from left to right across a period,
ionization energy increases.
 Within the same energy level electrons
experience an increasing pull from the
nucleus, so it takes more energy to remove
them.
Periodic Trends
 Moving down a group, ionization energy
decreases.
 The valence electrons feel less and less pull
from the nucleus as they get further from
the nucleus.
Periodic Trends
 2nd ionization energy is always greater than
the 1st ionization energy.
 When you remove an electron from an atom
the number of protons becomes greater
than the number of electrons. The
remaining valence electrons move closer to
the nucleus, making it harder to pull them
off the atom.
Periodic Trends
 As electrons are removed, ionization energy
increases gradually until an energy level is
empty, then it makes a big jump.
 Pulling an electron off of a alkali metal
(Group 1 elements) is easy. Trying to pull an
electron off of a noble gas (Group 18
elements) takes much more energy.
Periodic Trends
 Example: Put the following elements in
order of increasing ionization energy:
Sr, Cr, As, S, Rb, Cu
 Example: Put the following elements in
order of decreasing ionization energy:
O, V, K, P, Ga, Fr
The orbitals being filled for elements in
various parts of the periodic table.
Periodic Trends
 Which of the following elements will have a
very large second ionization energy? Third
ionization energy?
Na, Al, Ne, Mg, Si
The orbitals being filled for elements in
various parts of the periodic table.
Periodic Trends
 Ionic radius – similar to atomic radius but it
is the radius for an ion instead of an atom.
 Positive ions are always smaller than their
neutral atoms, and negative ions are always
larger than their neutral atoms.
 As you go down a group, ionic radius
increases.
Comparison of Atomic and Ionic
Radii
Periodic Trends
 As you go from left to right across a period,
positive ions decrease in size.
 Negative ions also decrease as you go across
a period, but they start off being much
larger than positive ions.
Periodic Trends
 Put the following ions in order of increasing
ionic radius:
 Hint: If all of the ions have the same
number of electrons, than the one with the
highest number of protons has the smallest
radius.
Na+1, Al+3, N-3, F-1, O-2, Mg+2
Periodic Trends
 Electronegativity – how strongly the nucleus
of an atom attracts the electrons of other
atoms in a bond.
 Nonmetals tend to gain electrons when they
form bonds, and have higher
electronegativities than metals, which tend
to lose electrons, when they form bonds.
Periodic Trends
 Moving from left to right across a period,
electronegativity increases.
 Moving down a group, electronegativity
decreases.
Electronegativities of the Elements
Periodic Trends
 Put the following elements in order of
increasing electronegativity:
Fe, Si, O, Ba, Ca, Cs
 Put the following elements in order of
decreasing electronegativity:
Se, F, Ag, Pt, Fr, Sb