Title: Electronic Structure and Periodic Properties
1Electronic Structure and Periodic Properties
- Wave Nature of Light
- Models of the Atom
- Bohr Model
- Quantum Mechanical Model
- Atomic Orbitals
- Electron Configurations
- Periodic Properties of Elements
2Electronic Structure of Atoms--Introduction
- Elements in the same group exhibit similar
chemical and physical properties. - Alkali Metals
- soft
- very reactive
- metal
- Noble Gases
- gases
- inert (unreactive)
- Why???
3Electronic Structure of Atoms--Introduction
- When atoms react, their electrons interact.
- The properties of elements depend on their
electronic structure. - the arrangement of electrons in an atom
- number of electrons
- distribution of electrons around the atom
- energies of the electrons
4Electronic Structure of Atoms--Introduction
- Understanding the nature of electrons and the
electronic structure of atoms is the key to
understanding the reactivity of elements and the
reactions they undergo. - Much of our knowledge of the electronic structure
of atoms came from studying the ways elements
absorb or emit light.
5The Wave Nature of Light
- Light is a type of electromagnetic radiation
- a form of energy with both electrical and
magnetic components
Wavelength (l) the distance between successive
peaks Frequency (u) the number of complete
wavelengths that pass a given point in 1 sec
6The Wave Nature of Light
The electromagnetic spectrum
7The Wave Nature of Light
- Different types of electromagnetic radiation have
different properties because they have different
u and l. - Gamma rays
- wavelength similar to diameter of atomic nuclei
- Hazardous
- Radio waves
- wavelength can be longer than a football field
8Quantized Energy and Photons
- Classical physics (mechanics) suggests that both
electromagnetic radiation and matter can have any
energy
A car rolling down a hill can have any potential
energy (energy of position) depending on its
position on the hill.
9Quantized Energy and Photons
- Classical mechanics is not correct, however.
- Max Planck suggested that energy is transferred
in packets called quanta (plural). - Quantum the smallest quantity of energy that
can be emitted or absorbed as electromagnetic
energy
10Quantized Energy and Photons
- Planck proposed that the energy of a single
quantum is directly proportional to its
frequency - E hu
- where E energy
- u frequency
- h Plancks constant (6.63x10-34 J-s)
11Quantized Energy and Photons
- According to Plancks theory, energy is always
emitted or absorbed in whole number multiples of
hu (i.e hu, 2hu, 3hu) - According to Plancks theory, the energy levels
that are allowed are quantized. - restricted to certain quantities or values
12Quantized Energy and Photons
- In order to understand quantized energy levels,
compare walking up (or down) a ramp versus
walking up (or down) stairs - Ramp continuous change in height
- Stairs quantized changed in height
- You can only stop on the stairs, not between them
13Quantized Energy and Photons
- If Plancks quantum theory is correct, why dont
we notice its effects in our daily lives? - Plancks constant is very small (6.63 x 10-34
J-s). - A quantum of energy (E hu) is very small.
- Gaining or losing such a small amount of energy
is - insignificant on macroscopic objects
- very significant on the atomic level
14Quantized Energy and Photons
- In 1905 Einstein used Plancks quantum theory to
explain the photoelectric effect. - Light shining on a clean metal surface causes the
surface to emit electrons. - The light must have a minimum
- frequency in order for electrons
- to be emitted.
15Quantized Energy and Photons
- Einstein explained these results by assuming that
the light striking the metal is a stream of tiny
energy packets of radiant energy (photons). - The energy of each photon is proportional to its
frequency. - E hu
16Quantized Energy and Photons
- When a photon strikes a metal surface
- Energy is transferred to the electrons in the
metal - If the energy is great enough, the electron can
overcome the attractive forces holding it to the
metal. - Any extra energy above the amount required to
free the electron simply increases the kinetic
energy of the electron.
17Quantized Energy and Photons
- Einsteins explanation of the photoelectric
effect led to a dilemma. - Is light a wave or does it consist of particles?
- Currently, light is considered to have both
wave-like and particle-like properties.
Matter also has this same dual nature.
18Models of Atomic Structure
- Scientists initially thought of the atom as a
microscopic solar system. - electrons orbiting the nucleus
- Unit 2 suggested that the atom has a tiny
positively charged nucleus with a diffuse cloud
of electrons surrounding it. - need better understanding of the nature of this
cloud of electrons.
19Atomic Models
- Two models are used to explain the behavior and
reactivity of atoms and ions. - Bohr model
- Quantum mechanical model
20Bohr Model
- Bohr developed an atomic model that explained the
line spectrum observed for the hydrogen atom.
- When an electrical current is passed thru a
sample of H2 (g), energy is transferred to the H2
molecules. - The molecules are broken up. The H atoms absorb
energy and jump to a higher energy level.
21The Bohr Model of the Atom
The H atoms relax back to their original energy
level by giving off the absorbed energy as
electromagnetic radiation.
High voltage
H2
22The Bohr Model of the Atom
The light is analyzed in a spectrometer by
separating it into its different colors.
High voltage
H2
23The Bohr Model of the Atom
The separated colors are recorded as spectral
lines.
High voltage
H2
Atomic spectrum
24The Bohr Model of the Atom
- The spectrum of atomic hydrogen consists of a
series of discrete lines such as the ones shown
previously. - Why would an atom emit only certain frequencies
of light and not all of them?
25The Bohr Model of the Atom
- According to the Bohr Model of the atom
- Electrons move in circular orbits around the
nucleus. - Energy is quantized
- only orbits of certain radii corresponding to
certain definite energies are allowed - an electron in a permitted orbit has a specific
energy (an allowed energy state)
26The Bohr Model of the Atom
- The allowed orbits have specific energies given
by the formula - En (-RH) 1 where n 1, 2, 3
- n2
- RH Rydberg constant 2.18 x 10-18 J
- n is called the principal quantum number
27The Bohr Model of the Atom
- Each orbit in an atom corresponds to a different
value of n. - As n increases the radius of the orbit increases
(i.e. the orbit and any electrons occupying it
are further from the nucleus) - n1 is the closest to the nucleus
- 0.529 Angstroms for the hydrogen atom
28The Bohr Model of the Atom
- The energy of the orbit is lowest for n1 and
increases with increasing n. - Lower energy more stable
- Lower energy more preferred state
29The Bohr Model of the Atom
- The lowest energy state of an atom is called the
ground state. - n 1 for the electron in a H atom
- When an electron has jumped to a higher energy
orbit (i.e. n 2, 3, 4) it is considered to be
in an excited state.
30The Bohr Model of the Atom
- To explain the line spectrum for hydrogen, Bohr
assumed that an electron can jump from one
allowed energy state to another. - Energy absorbed ? e- jumps to higher energy
state - e- relaxes back to a lower energy state ?
energy is emitted
31The Bohr Model of the Atom
32The Bohr Model of the Atom
- Since the energies of the orbits in an atom are
quantized, transitions from one allowed orbit to
another involves only specific amounts of energy. - DE Ef - Ei
33The Bohr Model of the Atom
- Since E hu, the energy of the light emitted can
have only specific values. - Therefore the u of the light can have only
specific values as well. - So, the line spectrum for each element will be
unique and will depend on the allowed energy
levels in that element.
34The Bohr Model of the Atom
- The Bohr model effectively explains the line
spectra of atoms and ions with a single electron - H, He, Li2
- Another model is needed to explain the reactivity
and behavior of more complex atoms or ions - Quantum mechanical model