Electrons in Atoms Chapter 5

Do you remember the 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 atom’s volume

Later Models Bohr – Planetary Model Schrodinger – Wave Mechanical Model

Problems Rutherford’s Model Didn’t Address 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?

Bohr Model Bohr - electrons in atom can have only specific amounts of energy NEW idea! each specific orbit is associated with specific amount energy 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 (shells) around the nucleus E3 n=3 n=3 E2 n=2 n=2 E1 n=1 n=1

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

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

energy levels get closer together the farther away they are from nucleus – not uniformly spaced 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 becomes excited & can move to higher energy level If electron loses specific amount of energy it drops down to lower energy level it gives off photon of light (color depends on wavelength of light given off)

When Matter is Heated it Gives Off Light

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

Example: fire works, pyrotechnics, flame test heat energy absorbed by the metal ions excites the atoms’ electrons absorbed energy is eventually released in the form of light

Example: light bulb electrical energy absorbed by the filament excites the atoms’ electrons absorbed energy is eventually released in the form of light

How this works: electrons absorb energy, get EXCITED, and “jump” to a higher energy level after a brief time, they “fall” back to a lower energy level, giving off a specific amount of energy (a quantum amount) in the form of a photon (colored light)

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

Absorption & Emission cannot easily detect absorption of energy by electron BUT can easily detect emission of energy by electron SEE: photons of light given off as excess energy is released

THE MYSTERY OF EMISSION SPECTRUMS there are two types: continuous spectrum bright line spectrum

Continuous Spectrum Solids, liquids, and dense gases emit light of all wavelengths, without any gaps

Bright Line Spectrum thin gases emit light of only a few wavelengths so see lines of color separated with gaps between them

atoms cannot emit energy continuously, rather they emit energy in precise quantities

Atomic Emission Spectra (AKA: bright line spectra) apply voltage across ends of glass tube containing gas - light is produced color of light depends on gas in tube every element produces its own unique color

our eyes see ONE color in the gas spectrum tube, however, if we use a prism we can see that each “color” is really multiple wavelengths of different colors

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

more examples: Hydrogen:

Practice Q

Which principal energy level of an atom contains electron(s) with the lowest energy? answer: a

What is total # of occupied principal energy levels in atom of neon in ground state? neon has 10 electrons: 1st shell: 2 2nd shell: 8 answer: b 1 2 3 4

nitrogen has 7 electrons: 1st shell: 2 2nd shell: 5 answer: a What is total # of fully occupied principal energy levels in atom of nitrogen in ground state? nitrogen has 7 electrons: 1st shell: 2 2nd shell: 5 answer: a 1 2 3 4

What is total # of electrons in completely filled fourth principal energy level? 8 10 18 32 2n2 2(42) = 32 answer: d

15 electrons total: phosphorus answer: d Which atom in ground state has five electrons in its outer level and 10 electrons in its kernel? 15 electrons total: phosphorus answer: d C Cl Si P

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

Li has 3 electrons answer: b Which electron configuration represents atom of Li in an excited state? 1-1 1-2 2-1 2-2 Li has 3 electrons answer: b

The characteristic bright-line spectrum of atom is produced by its electrons absorbing energy electrons emitting energy protons absorbing energy protons emitting energy answer: b