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
radial probability density
probability that we will find the electron at a specific distance from the nucleus. When plotted as a function of r, the distance from the nucleus, you get the radial probability function curve. The probability of finding the electron rises rapidly as we move away from the nucleus, maximizing at a distance of 0.529angstroms for an H atom, then falls of rapidly
p orbital
concentrated in two regions (lobes) on either side of the nucleus, separated by a node at the nucleus.
For each value of n, the three p orbitals for each shell have the same size and shape but differ in spatial orientation. The orientation is represented by the subscript px,py,pz (pyrepresents a p orbital along the y axis) As n increases, p orbitals increase in size
d orbitals
5 d orbitals:
dxy, dyz, dxz look like 4 leaf clovers and line on subscript plane with lobes in between axes
dx^2-y^2 looks like 4 leaf clover on xy plane with lobes along axes
dz^2 has two lobes along z axis and doughnut in xy plane.
all orbitals have the same energy
many electron atom vs H atom
In a H atom, the energy of an orbital depends only on its principal quantum number, n (all the subshells have same energy). In many-electron atoms, the electron-electron repulsions cause the different subshells to be at different energies
In a many-electron atom, for a given value of n, the energy of an orbital increases with increasing value of l. Energy increases in order of ns
George Uhlenbeck and Samuel Goudsmit
ELECTRONS BEHAVE IN PAIRS proposed electron spin: each electron behaves as if it were a tiny sphere spinning on its own axis.
Spin magnetic quantum number: ms, indicates the spin of electron, only has two values: +½ or -½. Spinning charge produces magnetic field. Opposite charges attract
pauli
EXCLUSION PRINCIPLE: no two electrons in an atom can have the same set of four quantum numbers n, l, ml, ms. Each orbital can hold a maximum of two electrons and they must have opposite spins
electron configuration
Electron configuration: way in which electrons are distributed among the various orbitals of an atom
Most stable electron configuration=ground state, electrons in lowest possible energy states. Orbitals are filled in order of increasing energy, with no more than two electrons per orbital
Electron configurations can be written by writing symbol for subshell and superscript to indicate the number of electrons in that subshell. Li (3 electrons)=1s22s1. Orbital diagrams depict orbitals by a box and each electron by a half arrow. An arrow pointing up represents positive spin and down=negative spin.
hund
OCCUPY ORBITALS SINGLY for degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized. Electrons will occupy orbitals singly to the maximum extent possible (parallel spins)
condensed electron configurations
using electron configuration of nearest noble-gas element of lower atomic number (noble-gas core) to abbreviate configuration. Na:[Ne]3s^1.
inner-shell electrons are called core electrons, outer shell electrons are valence electrons and are involved in chemical bonding
Elements in the same group (column) on PT have same valence electron configuration, so that is why they share many properties
transition metals electron configuration
The 4s orbital is filled before the 3d orbital (corresponds to transition metals) Mn: [Ar]4s^23d^5
rare earth elements electron configuration
The 14 elements corresponding to the filling of the 4f orbitals are lanthanide elements/rare earth elements. La: [Xe]6s^25d^1, Pr: [Xe]6s^24f^3
Actinide elements complete the 5f orbitals
Electron Configurations and the Periodic Table
Go from right→ left and top->bottom to know order in which orbitals are filled
The s block and p block of PT together are representative elements, sometimes called the main group elements
For representative elements we do not consider completely full d or f subshells to be among the valence electrons, and for transition elements we likewise do not consider a completely full f subshell to be among the valence electrons
abnormalities to electron configuraion
Transition metals in 4th column (group 6): nd^5ns^1instead of nd^4ns^2
Transition metals in 9th column (group 11):nd^10ns^1 instead of nd^9ns^1
