Download Atomic Emission Spectra, Electron Configuration, Periodicity

Atomic Emission Spectra,
Electron Configuration,
Periodicity
Ch.5
Coral Gables Senior High
Ms. Kiely
Pre-IB Chemistry I
Bell-Ringer #8
Today
1. Individual teacher-student meetings with rebuild
2. Unit 2: Ch.6 Introduction Questions DUE TODAY
1. State the group number(s) for each of the following groups of elements: the alkali metals,
the alkaline metals, the halogens, the noble gases, the transition metals. Take note that your
book refers to these groups as 1A-8A, and 3B-7B. TRANSLATE THESE INTO GROUPS 1-18!
2. Which group numbers correspond to the representative elements, and why are they given
that name?
3. Define atomic radius. What is the trend across a period for the atomic radius of an atom?
4. Define ionization energy. Does ionization energy pertain to cations or anions?
5. Define electronegativity. Explain why electronegativity increases left to right across a
period. Does electronegativity pertain to cations or anions?
Bell-Ringer #9
Turn in Unit 2 Introduction Questions
Bell-Ringer #9
Answer:
B
Bohr’s
Experiment
Bohr excited the atoms of hydrogen gas with an electric current (energy). He then
noticed a blue light emitting from the gas sample.
He had this blue light pass through a prism. The prism showed that the blue light was
composed of four specific wavelengths of visible light: red, blue-green, blue-violet, and
violet. He was observing hydrogen’s Atomic Emission Spectra
This is hydrogen’s Atomic Emission Spectra.
When the electrons of an atom absorb a particular amount of energy, a quanta, they
are able to “jump” to higher energy levels in the atom. We call this an electron’s excited
state.
The excited state is, however, unstable and the electron soon falls back to its original,
lower energy position in the atom, we call this its ground state. As it travels back
towards its ground state it releases the same amount of energy it originally absorb,
however it releases it in the form of electromagnetic radiation, or light.
Bohr’s Model of the Atom
This lead Bohr to propose that electrons travel
along well defined circular orbits, which he
referred to as energy levels, that are located
around the nucleus of an atom.
He figured since hydrogen always emitted the
same spectra of electromagnetic radiation, that it
was because electrons were located in specific
locations in the atom.
Although the Bohr model of the atom was able to
explain the emission spectrum of hydrogen with
great success, it fails to predict the emission
spectrum of atoms with more than one electron.
Why was Bohr’s model wrong?
1. Wave-Particle Duality: electrons actually behave as both particles and waves. Waves
are in continuous motion, and therefore are not localized to a specific position in
space.
2. Heisenberg’s Uncertainty Principle: we cannot know where an electron is at any
given moment in time - the best we can hope for is a probability picture of where the
electron is likely to be.
The possible positions of an
electron are spread out in
space in the same way that
a wave is spread across a
water surface.
Recall that the mere act of locate an object in space
(whether via naked eye or instrument) requires
light to reflect off of that object.
We cannot find an electron even if it is behaving as
a particle. The subatomic particles are so small
that light reflecting off of them will cause their
trajectory to change.
Why can’t we observe an electron directly?
Why do we believe electrons have particle-wave duality?
VIDEO: QUANTUM WORLD
Figure: beam of electrons going
through a slit produces a
diffraction pattern, a wave pattern,
instead of a straight line.
Evidence for particle-wave duality
of subatomic particles.
Schrödinger and the Electron Cloud
-Based on Heisenberg’s uncertainty principle, Erwin Schrödinger in 1926 developed
a mathematical equation to figure out where an electron might be located within
the electron cloud based on its level of energy.
-He knew that an electron’s location within the electron cloud could not be known
for certain- but a probability of where an electron might be could be calculated!
-We call areas of high probability of where an electron
might be an atomic orbital.
Atomic orbitals describe the three-dimensional areas where there is a high probability
that the electron will be located. They are regions around an atom’s nucleus in which
there is a 90% probability of finding the electron.
Shapes of orbitals will depend on the energy of the electron and its motion in that area.
When an electron is in an orbital of higher energy, it means it is most likely far from
the nucleus.
s orbitals are spherical
p orbitals are dumbbell-shaped.
There are three of them, located
on an x, y, and z axis.
Despite not being
able to directly detect
the location of an electron directly, scientists
Electron
Configuration
Diagram
like Schrödinger have provided mathematical formulas and theoretical models for us
that will help us determine where an electron MOST LIKELY is, as well as where an
electron may be traveling to when it absorbs energy and travels away from the
nucleus.
The formulas can be summarized in something called electron configuration.
1s_
We will fill in
diagonal arrows to
help us follow the
Aufbau Principle.
2s_
2p_ _ _
3s_
3p_ _ _
3d_ _ _ _ _
4s_
4p_ _ _
4d_ _ _ _ _
4f_ _ _ _ _ _ _
5s_
5p_ _ _
5d_ _ _ _ _
5f_ _ _ _ _ _ _
Electron Configuration Rules
1. Aufbau Principle: lowest energy orbitals are filled first
2. Pauli Exclusion Principle: up to two electrons can occupy an orbital, however
they must have opposite spins so as to reduce the repulsion of like charges. (Up
arrow means clockwise, down arrow means counter-clockwise.)
3. Hund’s Rule: electrons are placed into orbitals of a sublevel one electron at a
time, so as to minimize the amount of repulsion in one given orbital. Once all
orbitals in a sublevel are filled with one electron, then if more electrons remain
a second electron with opposite spin can be added to completely fill the orbital.
*Transition metals (metals located in groups 3-12) are exceptions to these rules.
To be discussed later.
1. Using the diagram, determine the electron configuration of the
following elements: Li, Mg, & Ar
2. State the full electron configuration of vanadium (V) and deduce the
number of unpaired electrons.
Answers:
1. Li: 1s²2s
Mg: 1s²2s²2p⁶3s²
Ar: 1s²2s²2p⁶3s²3p⁶
2. V: 1s²2s²2p⁶3s²3p⁶3d³4s²
There are three unpaired electrons.
Argon’s and Vanadium’s electron configurations give us a great example for a
different way of writing out an atom’s electron configuration, we call it
condensed electron configuration.
Ar: 1s²2s²2p⁶3s²3p⁶
V: 1s²2s²2p⁶3s²3p⁶3d³4s²
Notice that Vanadium’s electron configuration is similar to Argon’s since Argon
only has 18 electrons and Vanadium has 23. We can therefore condense or
abbreviate the electron configuration of Vanadium like this:
[Ar]3d³4s²
To use this form, you must abbreviate according to the noble gas that corresponds
to the specific element.
The electron configurations of the first 30 elements: