ch 6 electronic structure of atoms Flashcards
quantum theory
When lights are turned on, the electrons in atoms get excited to a higher energy by the electricity, and emit light when they drop back down to a lower energy
electromagnetic radiation
(aka radiant energy) carries energy through space. One form is visible light. All types of ER moves at speed of light (3.00*108m/s) and have wave like characteristics
wavelength/cycle
distance between two adjacent peaks/troughs of a wave
frequency
number of wavelengths that pass a given point each second. measured in hertz, Hz=s^-1, or /s
frequency-wavelenth formula
vlambda=c. v=nu, frequency (measured in cycles/sec=s-1= /s = Hz. lambda=lambda, wavelength (measured in meters). c= speed of light=3.00*108m/s
Long wavelength=lower frequency. Higher frequency=shorter wavelength
electromagnetic spectrum order12
(shortest wavelength, highest frequency) gamma rays, x rays, ultraviolet, visible light-(400-750nm)violet, blue, green, yellow, orange, red, infrared, microwaves, radio frequency (longest wavelength, lowest frequency)
wavelength units6
angstrom=10^-10m nm=10^-9m micrometer=10^-6(mew) mm=10^-3 cm=10^-2 m=1m
planck
QUANTIZED E
discovered: energy can be either released or absorbed by atoms only in discrete “chunks” of some minimum size. Quantum (fixed amount): smallest quantity of energy that can be emitted or absorbed as electromagnetic radiation. E of a single quantum equals Planck’s constant times frequency of radiation: E=h
h (Planck’s constant)=6.626*10-34joule-seconds (J-s). Matter is allowed to emit/absorb energy only in whole number multiples of h. 3h=three quanta of energy
einstein
PHOTOELECTRIC EFFECT emission of electrons from metal surfaces on which light shines. For each metal, there is a minimum frequency of light below which no electrons are emitted
Albert Einstein explains photoelectric effect: radiant energy behaves like a stream of energy packets, each energy packed called a photon. References Planck’s quantum theory to say energy of photon=E=h. Radiant energy itself is quantized
work function
A certain amount of energy (work function) is required for an electron to overcome the attractive forces that hold it in the metal. If photons hitting the metal have less energy than the work function, electrons do not escape metal surface. If photons have sufficient energy, electrons are emitted. If photons have more energy than work function, excess energy appears as kinetic energy of emitted electrons
monochromatic
radiation composed of a single wavelength
spectrum (continuous, line)3
Spectrum is produced when radiation from sources that produce different wavelengths is separated into its different wavelength components
Continuous spectrum:no blank spots in spectrum (ex: rainbow of colors of visible light)
Line spectrum: spectrum containing radiation of only specific wavelengths. Each wavelength represented by a colored line, which are separated by black regions (wavelengths that are absent from light)
rydberg’s equation
calculated the wavelengths of all the spectral lines of Hydrogen
Bohrs equation
E=(-hcRH)(1/n2)h=Planck’s constant, c=speed of light, Rh=rydberg constant, n=principal quantum number. -hcRH=-2.18*1018Jfor a H atom. This equation tells how much energy the electron will have, depending on which orbit it is in. n=1→ E is most negative. As n increases, E becomes less negative
bohrs model postulates3
ELECTRONS ONLY EXIST IN DISCRETE ENERGY LEVELS (QUANTIZED). ENERGY IS INVOLVED IN MOVING AN ELECTRON BETWEEN LEVELS
Only orbits of certain radii, corresponding to certain definite energies, are permitted for the electron
An electron in a permitted orbit has a specific energy and is in an “allowed” energy state. An electron in an allowed energy state will not radiate energy and therefore will not spiral into the nucleus
Energy is emitted or absorbed by the electron only as the electron changes from one allowed energy state to another. Energy is emitted/absorbed as a photon. e=h
principle quantum number n
corresponds to each orbit, Orbit closest to nucleus has n=1, radius of orbit gets larger as n increases.
Ground state of atom: lowest energy state (n=1)
Excited state: higher energy state of atom, n=2 or higher
Zero-energy state: When n gets infinitely large, E=0. Represents when electron is completely separated from nucleus. Zero-energy state is higher in energy than state with negative energies’
moving an electron between levels
Energy is absorbed for an electron to move to a higher state (n increases) and emitted when the electron jumps to a lower energy state (n decreases). changeE=Ef-Ei=Ephoton=hv. changeE=hv=hc/v=(-2.18*10-18J)(1/nf2-1/ni2). If nf
de broglie
MATTER WAVES
an electron when moving about the nucleus is associated with a particular wavelength: lambda=h/mv. h=planck’s constant m=mass (mass of electron=9.11*10-28g), v=velocity
The quantity mv for any object is its momentum
heisenberg
UNCERTAINTY PRINCIPLE
dual nature of matter places a fundamental limitation on how precisely we can know both the location and momentum of any object at the subatomic level
Heisenberg’s principle: the uncertainty principle: it is inherently impossible for us to know simultaneously both the exact momentum of the electron and its exact location in space. The more accurately one is known, the less accurately the other is known
Uncertainty principle: changex*change(mv)h/(4pi)
schrodinger
WAVE FUNCTIONS Represented by (psi). psi^2=probability density/electron density: represents the probability that the electron will be found at that location. If psi^2is large, there is high probability of finding an electron The wave functions in schrodinger’s equation are called orbitals. Each orbital describes a specific distribution of electron density in space, given by its probability density and each orbital has its characteristic energy and shape.
n quantum mechanical number
principal quantum number, n, can have positive integral values (1,2,3…). As n increases, orbital becomes larger and the electron spends more time farther from the nucleus. As n increases, electron has higher energy and is less tightly bound to nucleus. For H atom: En=-(2.18*1018J)(1/n2)as in Bohr model
collection of orbitals with same value of n are called electron shell. All the orbitals with n=3 are in the third shell.
shell with principal quantum number n will have n subshells
total number of orbitals in a shell is n^2
l quantum mechanical number
The azimuthal quantum number, l, can have values from 0 to n-1, for each value of n. Defines shape of orbital. l corresponds to letters : 0-s, 1-p, 2-d, 3-f
Orbitals with the same n andl are in the same subshell. Orbitals with n=3 and l=2 are called 3d orbitals and are in the 3d subshell
ml quantum mechanical number
The magnetic quantum number, ml, can have integral values between -l and l,including 0. Describes orientation of the orbital in space.
For each value of l, there are 2l+1values of ml
s orbital
spherically symmetric: electron density at a given distance from nucleus is same regardless of direction in which we proceed from nucleus
