Chapter 5 and 6 Test Flashcards
The Bohr Model
- Created by Neils Bohr
- Showed how energy of an atom changes when absorbing or emitting light
- Proposed that an electron is found only in specific circular paths, or orbits, around the nucleus.
- Each orbit= different energy level
- Absorbs energy–> higher orbit
- Emits energy–> lower orbit
Limitations of Rutherford’s Atomic Model
- RAM: Explained only a few simple properties of an element
- RAM: Could not explain the chemical properties of an element
Ground State
- When an electron is at the lowest energy orbit (stable)
- When at a high energy level, an electron is unstable)
What is the formula for energy emitted by an electron?
E=h x v
- V is the frequency
- H is a constant #
What are the two things that effect the color being emitted by an electron when releasing energy?
- The electrons in their outer energy levels are unique for each element
- The quantum of energy released. The quantum released is represented by the distance of energy levels .
Energy levels
- The fixed energies an electron can have
- Unequally spread upart
Quantum
-The amount of energy required to move an electron from one energy level to another energy level
The Bohr model vs. Rutherford’s Atomic Model
- RAM: could not explain why elements that have been heated to higher temps give off different colors of light
- Bohr model explains that when an atom emits light it is moving from one energy level to another
The Quantum Mechanical Model
- Created by Edward Schrodinger
- Devised a mathematical equation to describe the behavior of the electron in a hydrogen atom
- Came from mathematical solutions to the Schrodinger equation
The Quantum Mechanical Model vs. The Bohr Model
- The quantum mechanical model describes the energy of electrons like the Bohr Model
- Unlike the Bohr Model, QMM does not specify the exact path the electron takes around the nucleus
Orbitals
- A region of space an electron will be in at any given time
- Types: S,D, F, P
Sublevel
-The shape of an orbital
Aufbau principle
-Electrons occupy the orbitals of the lowest energy level first
Hund’s rule
- Every orbital in a subshell gets one electron before any orbital gets 2 electrons
- Single electrons have parallel spin
Pauli Exclusion Principle
-To occupy the same orbital, two electrons must have opposite spins.
Three ways to show electron configurations
- Aufbau Diagram
- Standard Electron Configuration
- Orbital filling Diagram
Amplitude
-The wave’s height from zero to the crest
Wave length
- The distance between the crests
- Represented by the greek letter lamba (upside down y)
Frequency
- The number of wave cycles to pass a given point per unit of time
- Represented by the greek letter nu (v)
The Frequency and Wave Length of Light
- Inversely proportional ( As one decreases, the other increases)
- As wave length increases, frequency decreases
Electron Magnetic Radiation
- Massless yet has energy
- Light
- Examples: radio waves, visible light, X-rays, gamma rays
- Higher the frequency, the more dangerous it is
Atomic Emission Spectra
AES: The pattern formed when light passes through a prism or diffraction grating to separate it into the different frequencies of light it contains
- Ex: White light separates to the colors of the rainbow
- Unique finger print of an element (each 1 is different)
Spectra Line
-the line that forms when white light is separated using the Atomic Emission Spectra
How energy travels from one place to another
- By traveling on a particle (particle is charged then moved)
- By traveling as a wave
Energy traveling by particle
- Follow classical mechanics (laws of motion, projectile motion)
- Have energy and mass
Energy in wave motions
- Electromagnetic radiation travels as waves
- Are massless and have energy
- Travels in chunks (photons)
- Wave Particle Duality: the characteristics of light, it travels like a wave behaves like a particle
Photons
- Chunks or quanta of energy that behave as though they were particles
- 1 quanta= 1 photon
- Energy of photon= h x v
Planck’s constant
- (h in E=h x v)
- 6.626 x 10^-32 j x s
Why the Atomic Emission Spectra is Unique
- Every element has a unique number of electrons which occupy unique orbitals and energy levels
- Because the electrons and the orbitals/energy levels they occupy is unique, the amount of energy released is unique.
- Energy released (photons) effects frequency and wavelength directly
Dmitri Mendeleev
- Constructed the periodic table of elements
- Arranged the elements in his periodic table in order of increasing atomic mass
What are the 3 classes of elements?
- Metals
- Nonmetals
- Metalloids (characteristics of metals and nonmetals)
Metals
- 80% of periodic table
- Good conductors of heat and electricity (best= silver, second best= copper)
- Freshly cleaned metal will have a sheen (can reflect light)
- All metals are a solid @ room temp. (except mercury)
- Metals= ductile (can be turned into wire)
- Most metals= malleable ( able to be made into sheets without breaking)
Nonmetals
- Most nonmetals are gases @ room temp
- few= solids (phosphorus, sulfur)
- Nonmetal that is liquid= bromine
- Because of many different types of nonmetals, no set characteristics
- Hard non-metals (diamond), brittle (break if hit)nonmetals (phosphorus)
- Poor conductors of heat & electricity
Metalloids
- Has properties similar to metals and nonmetals
- Under some conditions, act like metals, in other conditions act like nonmetals
- Poor conductor of electricity
How the Atomic Emission Spectra works with the Bohr Model
- When an electron goes from a higher energy level to a lower energy level, or vice verse, it emits or absorbs energy
- Depending on the amount of quantums/ energy levels moved down or up, an amount of energy will be absorbed or given off
- Higher the energy, higher the frequency
- Frequencies of light give off a color depending on the frequency (red to purple, lower to higher frequency)
- Spectra lines on the AES depend on how much electrons and element has and how much energy the electrons are giving off
- E=hv means higher energy of the light has a higher frequency
