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Electrons in Atoms
Models of the Atom
 The chemical properties of atoms, ions and
molecules are related to the arrangement of the
electrons within them.
 J.J. Thomson came up with the plum pudding
(chocolate chip cookie dough) model to explain
what he thought the atom looked like.
 Ernest Rutherford said electrons surround the
nucleus of the atom.
Ernest Rutherford’s Findings
 Rutherford said electrons surround the nucleus
of the atom and that:
 Atomic nucleus is composed of protons
(positive) and neutrons (neutral).
 The nucleus is very small compared to the
size of the entire atom.
 Questions left unanswered:
 How are electrons arranged and how do they
move?
The
Rutherford
atom.
Models of the Atom
 Niels Bohr said that electrons are arranged around
the nucleus in paths or orbits.
 Bohr thought that electrons were in fixed paths
around the nucleus and that they could not lose
energy and fall into the nucleus.
Electromagnetic Radiation
 Classical physics says matter is made up of
particles, and that energy travels in waves
 Electromagnetic Radiation is radiant energy,
both visible and invisible
 Electromagnetic radiation travels in waves
 Electromagnetic radiation given off by atoms
when they have been excited by any form of
energy
 flame tests
 All waves are characterized by their velocity,
wavelength, amplitude, and the number of waves
that pass a point in a given time
Physics and the Quantum Mechanical
Model
Light and Atomic Spectra
 Isaac Newton thought light consisted of particles.
 By 1900, most scientists thought that light was a
wave phenomenon, and consisted of
electromagnetic waves.
Physics and the Quantum Mechanical
Model
 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 light particle with a
wavelength of 675 nm?
 What is the energy of a photon with a frequency of
2.94 x 1015 cycles per second (s-1 or Hz)?
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
 Basis of quantum theory
 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 quanta. The energy of photons is
quantized according to the equation E = hu.
 Light was therefore thought to have a dual waveparticle 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.
Problems with Rutherford’s
Nuclear Model of the Atom
 Electrons are moving charged particles
 Moving charged particles give off energy
 Therefore the atom should constantly be giving off
energy
 And the electrons should crash into the nucleus
and the atom collapse!!
Bohr’s Model
 Explained spectra of hydrogen
 Energy of atom 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
Niels Hendrik
David Bohr
(1885-1962)
Source:
Emilio Segre
Visual Archives
Bohr’s Model
 Atoms have a minimum energy called the ground
state
 therefore they do not crash into the nucleus
 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
Problems with the Bohr Model
 Only explains hydrogen atom spectrum
 and other 1 electron systems
 Neglects interactions between electrons
 Assumes circular or elliptical orbits for electrons which is not true
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 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 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 per orbital.
 The maximum number of electrons that can
occupy a principal energy level is given by the
formula 2n2, where n is the principal quantum
number.
 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.
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
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
 Gallium
 Thorium
 Platinum
Noble Configurations
 What element is represented by the following noble
gas configuration:
 [Kr]5s24d105p2
 [Ar]4s2
 [Xe]6s24f145d6
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 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
The Modern Periodic Table
 Columns are called Groups or Families
 Rows are called Periods
 Each period shows the pattern of properties repeated
in the next period
 Main Groups = Representative Elements
 Transition Elements
 Bottom rows = Lanthanides and Actinides
 really belong in Period 6 & 7
Metallic Character
 Metals
• Metalloids
 malleable & ductile
 shiny, lustrous
 conduct heat and
electricity
 most oxides basic
and ionic
 form cations in
solution
 lose electrons in
reactions - oxidized
• Nonmetals
 Also known as
semi-metals
 Show some
metal and
some nonmetal
properties
 brittle in solid state
 dull
 electrical and
thermal insulators
 most oxides are
acidic and molecular
 form anions and
polyatomic anions
 gain electrons in
reactions - reduced
Metallic Character
 Metals are found on the left of the table,
nonmetals on the right, and metalloids in
between
 Most metallic element always to the left of the
Period, least metallic to the right, and 1 or 2
metalloids are in the middle
 Most metallic element always at the bottom of a
column, least metallic on the top, and 1 or 2
metalloids are in the middle of columns 4A, 5A,
and 6A
Reactivity
 Reactivity of metals increases to the left on the Period
and down in the column
 follows ease of losing an electron
 Reactivity of nonmetals (excluding the noble gases)
increases to the right on the Period and up in the
column
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
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
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.
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.
Comparison of Atomic and Ionic
Radii
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