chapter 5 - electrons in atoms

Neil Bohr

• Was Rutherford’s student and added to his model → how energy changes when it absorbs or emits light 

• Proposed that an electron is found only in fixed paths or orbits around the nucleus 

• Each electron orbit has a fixed energy, which is called energy levels 

• To move energy levels an electron must gain or lose just the right amount of energy (quantum)

  • Bohr's model failed to explain energy is absorbed and emitted by atoms with more than one electron bc he used hydrogen as the example 

Rutherfords limitations

• Didn't explain the chemical properties of elements

  • why metals/compounds of metals give off characteristic colors when heated

  • Needed a model that better described the behavior of electrons in atoms to explain what led to the chemical properties of elements 

Erwin Schrodinger

Found the ‘Quantum Mechanical Model’ which determines the allowed energy an electron can have and how likely it is to find the electron in various locations around the nucleus

• Electron cloud 

  • More dense = higher probability  → less dense = lower probability

  • Nucleus represented as a fuzzy cloud

  • No boundary to the cloud bc a slight chance of finding electrons at a considerable distance away from the nucleus

    • Attempts to show probability as a fuzzy cloud limited to volume in which electron is found 90% of the time 

• Sim to the Bohr model

  • Restricts the energy of electrons to certain values 

• Diff to Bohr model

  • Doesn't specify an exact path the electron takes around the nucleus

Max Plank

Wanted to explain why a body first appears black, then red, then yellow, then white as its temperature increases 

Found that he could explain color changes if he assumed that the energy of a body changes only in small units, or quanta 

• Showed mathematically that the amount of radiant energy (E) of a quantum absorbed/emitted by a body is proportional to the frequency of radiation (v) 

  • Small energy change means low frequency

  • Big energy change means high-frequency

• Also found that the amount of energy of one quantum equals the frequency x a constant  

• Planck’s constant 

  • 6.626 times 10^-34 joules 

Quantum

• Electrons cant move in between energy levels 

• A quantum of energy equals the amount of energy needed for electrons to move from one energy level to another

  • The amount of energy is not always the same 

  • Higher energy levels are closer together and far from the nucleus so they take less energy

  • Lower energy levels take more energy bc closer to the nucleus and far away from each other  

 E = hv

energy=planck’s constant * frequency 

electron configuration

Most stable arrangement of electrons in an atom → unexcited state 

electron configuration rules

1. Aufbau Principle 

   • Electrons fill orbitals of lower energy levels first 

     • Look at the models in notes

   • Boxes can overlap in diagrams

     • Based on how much energy a level has 

2. Pauli Exclusion Principle 

   • Orbitals can occupy at most 2 electrons

     • To occupy the same orbital, two electrons must have opposite spins 

     • Think of spin as counterclockwise and clockwise 

3. Spin 

   • Electrons in the same orbital must be spin-paired (one up one down) 

4. Hund’s Rule 

   • Electrons fill orbitals of sublevels in a way that maximizes the spin in one direction 

     • It would go  ^ ^ ^ instead of ^>^

frequency

how many waves pass through a given point per unit of time S^-1→ cycles per second) 

• units = hertes

wavelength

distance from crest to crest

• unit=meters

relationship between frequency and wavelength

they are inversely proportional

• higher frequency = smaller wavelength —> lower frequency = bigger wavelength

Electromagnetic radiation (waves) travel at the speed of light

Scale ( lower to higher frequency - bigger to smaller wavelength)

1. Radiowaves 

2. Radar 

3. Microwaves 

4. Infrared 

5. Visible light spectrum

   1.  Red, orange yellow, green, blue, violet 

6. Ultraviolet (UV) 

7. X-rays 

8. Gamma rays 

Colors in contrast to wavelength

Color is determined by wavelength range and energy needed 

• Red has the least amount of energy 

  • Longest wavelength

• Violet has the most 

  • Shortest wavelength

• The longer the wavelength = lower the energy → shorter wavelength = higher the energy

Color to wavelength:  

Red = 625-740

Orange = 590-625

Yellow = 565-590

Greens = 520 - 565

Blues = 440 - 520

Violets = 380 - 440

frequency and wavelength equation

C = wavelength*frequency

atomic orbitals

regions of space where there is a high probability of finding an electron

• same name as sublevels

• can hold up to 2 electrons on each sublevel

• come in different sizes and shapes

sublevels

s, p, d or f

1. Each energy level has a specific # of sublevels

2. The number of sublevels=number of that energy level, for example: 

1^st energy level has 1 sublevel - 1s

2^nd energy level has 2 sublevels- 2s and 2p

3^rd energy level has 3 sublevels- 3s, 3p and 3d

4^th energy level has 4 sublevels- 4s, 4p, 4d and 4f

3. Sublevel energies from lowest to highest: s,p, d, f

4. A specific number of electrons can go into each sublevel:

1. s-max of 2 e-

2. p-max of 6 e-

3. d-max of 10 e-

4. f-max of 14 e-

planks constant

6.626×10^-34 joules

all equations needed

1. △E=hc/λ

2. E=hv

3. C=λv

C= speed of light

v = frequency

λ = wavelength

h = planks constant

△E = difference between two energy levels

speed of light

3.0×10^8 m/s