SECTION 1 - QUANTUM MECHANICAL MODEL OF ATOM Flashcards
Quantum mechanics
Model that describes how electrons exist in atoms and helps explain chemical and physical properties of the elements
Parts of an atom:
- Proton (p)
- Neutron (n)
- Electron (e-)
Electromagnetic radiation
Tool to understand electronic structure in atom and consists of a wave of energy
Frequency
Number of cycles per second (s-1) (v)
Wavelength
Distance the waves travels during one cycle (nm, m , um) (lamda)
Amplitude
Height of crest of each wave
Relationship between frequency and wavelength
Inversely proportional. Greater frequency = small wavelength, small frequency = big wavelength)
What type of ER has the longest wavelength (smallest frequency)
Radio
What type of visible light has the shortest frequency
Red
What type of ER has the shortest wavelength (biggest frequency)
Gamma ray (produced by sun, stars, etc)
Speed of light
2.998 x 10^8 ms-1
Relationship between visible light and brightness/color
Amplitude determines the brightness and frequency/wavelength determines color
Interference
Phenomenon in which 2 waves combine by adding or subtracting waves
In phase (constructive interference)
Waves are in line and combine together = higher amplitude, higher intensity (bright)
Out phase (destructive interference)
Waves are not in line and subtract one another = no amplitude, no intensity (dark)
Wave-particle duality of light
Waves can exhibit particle like properties and particles can exhibit wave-like properties (photons)
Photoelectric effect
Albert Einstein demonstrated that light has properties of a particle. The experiment consisted of the emission of electrons when electromagnetic radiation hits a metal
2 reasons why light is not proven as a wave (photoelectric effect) (2)
- Light had to be a certain minimum frequency or no electron was ejected (if light was a wave, any frequency of light would work)
- There was no lag time and the current flowed as soon as light was hit at the metal (if light was a wave, dim light would have taken a while to have enough energy to kick out the electron)
Photon
Term by Einstein of a particle of light and viewing it as a “packet” of electromagnetic radiation and has characteristics of Plank’s constant
Planck’s constant
6.626 x 10^-34
Blackbody radiation
Conducted by Max Planck. Blackbody is an idealized object that absorbs all electromagnetic radiation it comes into contact with. It then emits thermal radiation in continuous spectrum according to temperature. Planck theorized that energy emitted by hot object is in the form of packets (photon)
Ephoton formula
h = c/lamda
What is the energy of a packet dependant on
Dependant on frequency of light. Photon of the right frequency hits an electron on the surface and supplies the right amount of kinetic energy to kick electron
Kinetic energy formula
KE = energy of photon - binding energy
Atomic spectra
The spectrum of the electromagnetic radiation emitted or absorbed by an electron during transitions between different energy levels within an atom. Viewable via prism
Niels Bohr
Developed a model for the atom to explain the atomic line spectra for the H atom
What did Bohr popstulate (3)
- Electrons can move in a circular orbit (wrong)
- Electrons have a fixed set of allowed orbits (energy states) at fixed distances from nucleus with principle quantum number n (correct)
- Electrons can jump from one allowed orbit to another with a certain energy
Energy of electron in an orbit formula
En = -2.179 x 10^-18J (z^2/n^2)
Excitation
Electrons absorbs energy and gains enough energy to move to a higher orbit
Relaxation
Electron releases energy (emitted) to go to a lower orbit
Ground state
Lowest energy level (n=1)
Excited state
Other energy levels (n = or 2 or greater)
Difference in energy (delta E = Ef - Ei) formula
(-2.179 x 10^-18 J)(z^2/n^2f - z^2/n^2i)
Ionization energy
Minimum amount of energy needed to remove 1 electron from a gaseous ground state atom to an infinite distance
Louis De Broglie
Suggested that small particles like electrons have wave length properties (contrary to Einstein)
What is the wavelength dependant on
Mass and velocity
De Broglie relation formula
lamda = h/m x v
Mass of electrons
9.1094 x 10^-31 Kg
Heisenburg uncertainty principle
States that we cannot know both the position and speed of a particle (photon/energy) with perfect accuracy
Heisenburg uncertainty principle formula
Delta x = m delta v -> h/4pi
Relationship between position/velocity and uncertainty
Inversely proportional (the more uncertain in position, the smaller uncertainty in velocity)
Schroiders equation
Describes the probability of finding an electron in an atom
Dual nature of electrons
Electrons in hydrogen atoms are seen as standing waves due to their dual nature (similar to light)
Consequences of Schroiders equation (3)
- Electron energy is quantized (specific allowed values)
- Each energy level has 1 or more associated wave functions
- Squaring wave function gives high probability map of electron position in 3 dimensional space (orbital)
3 quantum numbers:
- Principle quantum number (n)
- Angular momentum quantum number (l)
- Magnetic quantum number (mI)
Principle quantum number (4)
- Indicates size and orbital energy
- Integer
- Larger n = larger orbital
- As n increases, so does the distance from nucleus and energy increase
Angular momentum quantum number (3)
- Indicates orbital space
- Integer
- Every number from 0 up to n-1
Magnetic quantum number (3)
- Determines orientation of orbital in space
- Integer from -l through 0 to +l
- The number of ml = number of orbitals
How are shapes of orbitals calculated
Based on the probability of finding an electron in a specific location of space outside nucleus 90% of the time
S orbital
Spherical shape and size increases as n increase
Radial/spherical nodes
Circular regions of 0 probability of finding electrons in 2s and 3s orbitals (# of radial nodes = n-l-1)
P orbital
First occurs in level n=2 and has 3 different orientations that are the same size, shape, energy
What is found at the nucleus of every p orbital
Angular node
Spin orientation quantum number (ms)
Electron spin orientation and is +1/2 if arrow is up and -1/2 if arrow is down
Degenerate
Orbitals with the same energy
Effective nuclear charge Zeff:
Actual amount of positive charge experienced by an electron in a multi electron atom (Zeff = Z - S). Z is atomic number, S is core electrons
3 rules for electron configuration:
- Aufbau principle
- Pauli exclusion principle
- Hunds rule
Aufbau principle
Orbitals of lowest energy filled first
Pauli exclusion principle
Each electron in atom has a unique set of 4 quantum numbers meaning only 2 electrons placed in orbital must have opposing spins
Hunds rule
Electrons singally filled in each orbital of same energy before pairing (easier to decrease repulsion)
Valence electrons
Electrons in the largest and least stable occupied orbitals (outer electrons)
Core electrons
lectrons in the smaller and more stable occupied orbitals (inner electrons). Core electrons have same configuration as preceding noble gas in table
Short hand notation
Using noble gas to represent core electron configuration with valence electrons (eg. [He])
Important exceptions for transition metals
Transition metals have general configuration of 4s2 3d# but electrons can be removed from 4s to satisfy full shell of d
Periodic table blocks (4)
- S block: alkali metal + alkali earth
- P block: basic metal + metalloid + nonmetal + halogen + noble gas
- D block: transition metal
- F block: lanthanoids + actinoids
Diamagnetic atoms
Atom/ion with all electrons paired and has no magnetic properties
Paramagnetic atoms
Atom/ion with unpaired electrons and shows net magnetic properties. The amount of paramagnetism = number of unpaired electrons
Main group metals tend to ___ enough electrons to adopt the previous noble gas configuration
Lose
Main group nonmetals tend to ___ enough electrons to adopt the next noble gas configuration
Gain
Transitional metal ion configurations
Transition metals first lose ns electron before d electrons to form cations
How do we perceive color
Through the presence of a variety of wavelengths in white line is how we perceive color in objects
X ray
Passes through substances, then blocks visible light and used to image bones and internal organs. Can cause cancer
UV
In between x ray and visible light and is the component of sunlight that produces burns or tans. Can cause cancer
Visible light
Ranges from violet (shortest wavelength, higher energy) to red (longest wavelength, lower energy). Impacts certain molecules in our eyes to change shape and sends signal to our brains that results in our ability to see
Infrared radiation
The heat that is felt when hand is placed near hot object
Microwave
Efficiently absorbed by water and can heat substances that contain water
Radio waves
Transmits signals responsible for AM and FM radio, telephone, tv
