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: nucleus very small: all the positive charge & most of mass of atom nucleus very small: only 1/10,000th 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 Bohr - electrons in atom can have only specific amounts of energy NEW idea! 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. E3 n=3 n=3 E2 n=2 n=2 E1 n=1 n=1

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

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

Max Capacity of Bohr Orbits 2n2 n 32 4 18 3 8 2 1 Max # of Electrons Orbit

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 in energy between 2 levels 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? Potential Energy

Ground State vs. Excited 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?

What is total # of occupied principal energy levels in atom of neon in ground state? 1 2 3 4

What is total # of fully occupied principal energy levels in atom of nitrogen in ground state? 1 2 3 4

What is total # of electrons in completely filled fourth principal energy level? 8 10 18 32

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

Which electron configuration represents atom in excited state? 2-8-1 2-8 2-7-1

Which electron configuration represents atom of Li in an excited state? 1-1 1-2 2-1 2-2

The characteristic bright-line spectrum of atom is produced by its Electrons absorbing energy Electrons emitting energy Protons absorbing energy Protons emitting energy