Electrons in Atoms

Chapter 5

What were early steps in development of atomic theory?

• John Dalton – Billiard Ball Theory – Atom was indivisible • J.J. Thomson – Plum Pudding Model – Atom was composed of smaller particles

Rutherford Model

• nucleus contains: – all the positive charge & most of mass of atom • nucleus very small: – only 1/10,000 th of atomic diameter • electrons occupy most of volume

Later Models

• Bohr – Planetary Model • Schrodinger – Wave Mechanical Model

Problems with the Rutherford Model • Why don ’ t electrons crash into nucleus?

• How are electrons arranged?

• Why do different elements exhibit different chemical behavior?

• How is atomic emission spectra produced?

Atomic Emission Spectra

• gas in glass tube & apply voltage across ends – produces light • color of light depends on gas in tube • every element produces its own unique color

emission spectrum of element is set of frequencies (or wavelengths) emitted

Why is emission spectra useful?

• use it to determine if given element is present in sample • Neon lights

Emission & Absorption Spectra of Elements

Bohr Model

• each specific amount energy is associated with specific orbit – electrons restricted to these orbits • Bohr assigned quantum number (n) to each orbit – the smallest orbit (n= 1) • closest to nucleus • has lowest energy – larger the orbit, more energy it has

Bohr Diagram

• Shows all the electrons in orbits or shells about the nucleus.

n=3 E 3 n=3 n=2 E 2 n=2 n=1 E 1 n=1

Bohr Model

• energy absorbed when electron: – moves to higher orbit – endothermic process (farther from nucleus) • energy released when electron: – drops to lower orbit – exothermic process (closer to nucleus)

energy levels get closer together the farther away they are from nucleus Larger orbits can hold more electrons

Max Capacity of Bohr Orbits

Orbit Max # of Electrons 4 n 1 2 3 2 8 18 32 2n 2

Electron Transitions

• If electron gains (absorbs) specific amount of energy – it can be excited to move to higher energy level • If electron loses specific amount of energy – it drops down to lower energy level

Hydrogen has 1 electron, but it can make many possible electron transitions

Absorption & Emission

• cannot easily detect absorption of energy by electron BUT • can easily detect emission of energy by electron – photons (light) given off as excess energy is released

Emitted Light

• energy of emitted light (E = h  – matches difference levels in energy between 2 • don ’ t know absolute energy of energy levels, but – observe light emitted due to energy changes

ladder often used as analogy for energy levels of atom How is this one different?

Ground State vs. Excited State

• Ground state: – lowest energy state of atom – electrons in lowest possible energy levels • configurations in Reference Tables are ground state • Excited state: – many possible excited states for each atom – one or more electrons excited to higher energy level

Success of Bohr

’

s Model

• Bohr ’ s model could predict frequencies in emission spectrum of hydrogen • Predicted correct size of H atom • Unfortunately, didn ’ t work for anything with more than 1 electron

Which principal energy level of an atom contains electron with the lowest energy?

a) n=1 b) n=2 c) n=3 d) n=4

What is total # of occupied principal energy levels in atom of neon in ground state?

a) 1 b) 2 c) 3 d) 4

What is total # of fully occupied principal energy levels in atom of nitrogen in ground state?

a) 1 b) 2 c) 3 d) 4

What is total # of electrons in completely filled fourth principal energy level?

a) 8 b) 10 c) 18 d) 32

Which atom in ground state has five electrons in its outer level and 10 electrons in its kernel?

a) C b) Cl c) Si d) P

Which electron configuration represents atom in excited state?

a) 2-8-2 a) 2-8-1 b) 2-8 c) 2-7-1

Which electron configuration represents atom of Li in an excited state?

a) 1-1 b) 1-2 c) 2-1 d) 2-2

The characteristic bright-line spectrum of atom is produced by its a) Electrons absorbing energy b) Electrons emitting energy c) Protons absorbing energy d) Protons emitting energy