Chapter 7 Atomic Structure

Chapter 7Atomic Structure

Electromagnetic Radiation • Light is a form of electromagnetic (EM) radiation • All forms of EM radiation are types of kinetic energy • Page 239

EM Radiation • Describe each form of EM radiation by its: • Wavelength • Frequency • Energy of its photon

All forms of EM radiation travel at the speed of light (c) Wavelength x frequency = speed of light

EM Radiation • The longer the wavelength the: • lower the frequency and the lower the energy of the EM radiation • The shorter the wavelength the: • higher the frequency and the higher the energy of the EM radiation

Planck’s Constant • Energy of a photon … • Pages 240/241

EM Radiation • Compare the wavelength, frequency, and energy of: • Ultraviolet light and infrared light

EM Radiation • The higher the energy of the EM radiation the more damaging it is to living tissue. • CONSIDER GAMMA RAYS, X-RAYS, AND UV LIGHT

Emission of Energy by Atoms • Thanks to the work of Planck and Bohr we know that: • When atoms are energized by an input of energy their electrons are excited (energized) • When excited electrons return to lower energy states they emits energy in the form of light. • Emits photons of energy • Energy of the photons emitted depends upon how excited the electron was.

Photoelectric Effect • Electrons will leave a metal when light of sufficient frequency strikes the metal. • Called the threshold frequency • Light < threshold frequency – no electrons emmitted • Page 244

Photoelectric Effect • Light > threshold frequency the • Number of electrons increases with intensity of the light • Kinetic energy of the electrons emitted increases with the frequency of the light • Lead Einstein to…. E = mc2

Planck and Einstein • Energy is quantized • Occurs in discrete packets called quanta • Similar to $ comes in discrete quantities, penny, nickel, dime, quarter • EM radiation has both wave properties and particle/matter properties • Each wavelength is associated with a specific quantum of energy

Bohr Experiment (~1911) • Bohr excited hydrogen atoms by running electricity through a tube of hydrogen gas. • The gas gave off a pink light.

Bohr Experiment (~1911) • Bohr aimed a beam of the pink light at a “prism” • Found the pink light generated a line spectrum not a continuous spectrum • Line spectrum – specific colors of light observed • Continuous spectrum – all colors of light present

Bohr Experiment (~1911) • He observed 4 bands of color: (pg 247) • Purple (410 nm) • Blue (434 nm) • Green (486 nm) • Red (656 nm)

Bohr Experiment (~1911) • He then calculated the energy of each color of light • _____________ was the highest energy and _____________ was the lowest energy.

Bohr’s Interpretation of the Data • Bohr proposed: • The electrons circle the nucleus in orbits of specific energies. • Electrons are always in one of the circular orbits. • Larger orbits are of higher energy than smaller orbits

The electricity excites electrons and allows them to move to higher energy orbits. • When the excited electrons return to lower energy orbits they emit energy in the form of light. • For each orbit change, the difference in energy between the orbits corresponds to the energy of the light emitted. • Bohr concluded that because 4 specific wavelengths of light are emitted by hydrogen there are 4 possible orbit changes

Bohr Model of the Atom

Bohr’s Model • When Bohr’s mathematical approach was applied to other elements it didn’t work. • Bohr’s model of the atom has been revised to replace the circular orbits with “wave mechanical model” of the atom

Modern Atomic Structure • Still picture electrons to be at specific energy levels, but no longer picture them as traveling in circular orbits. • The current model of the atom locates electrons in orbitals.

Orbitals • Each orbital is of a specific energy, size, and shape • Each orbital can hold a maximum of 2 electrons of opposite spin (Pauli exclusion principle)

Orbitals • The exact path of an electron in an orbital is not known. • Heisenberg uncertainty principle states that it is impossible to determine the location and path of an electron at the same time

Orbitals • Orbital shapes describe the region in space where an electron will be found 90% of the time. • Each orbital is described by 3 quantum numbers….and each electron by 4 quantum numbers

Modern Atomic Theory • Atoms have specific energy levels in which electrons may be found. • Called Principal Energy Levels (PEL) • PEL farther from the nucleus are larger and of higher energy. • Assign a number (n) to each PEL • See board

Modern Atomic Theory • Within each PEL are sublevels • Sublevels are named: s, p, d, and f • The larger the PEL the more sublevels it contains

Modern Atomic Theory

Modern Atomic Theory • Sublevels contain orbitals.

Describing Orbitals • See pages 257/258 for diagrams of the orbitals • S orbitals are spherical • The 3 p orbitals are shaped

Putting it All Together • Unless they are excited, electrons always occupy the lowest energy orbital with room. • Electrons enter orbitals of a given sublevel one at a time before pairing up (Hund’s rule) • Consider 2 electrons in a p sublevel:

The fun part! • Our goal is to write the following for atoms and ions: • Electron configuration • Box/energy diagram • Lewis dot symbol • Our goal is also to: • Identify core and valence electrons

Terms • Electron configuration – shows the number of electrons in each sublevel • Box/energy diagram – shows the number of electrons in each orbital • Orbitals are shown as boxes • electrons are shown as arrows

Terms • Lewis Dot Symbol – shows the valence electrons as dots around the symbol for the element • Maximum of 2 electrons per side of the symbol

Terms • Valence electrons – all the electrons in the highest occupied PEL • Valence electrons are the ones involved in bonding • Core electrons – all electrons not considered valence electrons

Still to Come in January • Calculations in 7.4 • 7.5 – 7.10 in detail • 7.12

Chapter 7 Atomic Structure