×
Log in to StudySoup
Get Full Access to UT - Class Notes - Week 2
Join StudySoup for FREE
Get Full Access to UT - Class Notes - Week 2

Already have an account? Login here
×
Reset your password

UT / Chemistry / CHEM 120 / Which unit is used to measure wavelength?

Which unit is used to measure wavelength?

Which unit is used to measure wavelength?

Description

Week 2 Chem Notes: Chapter 2 09/09/2019


Which unit is used to measure wavelength?



● The Nature of Light: It’s Wave Nature

○ Light: a form of electromagnetic radiation

■ Light is composed of perpendicular oscillating, or swinging, waves - one for electric field, one for magnetic field 

■ Light is one of the ways by which energy travels through space

● Characteristics of a Wave

○ Wavelength(λ): distance between two peaks

○ Frequency(ν): number of waves, or cycles, per second that pass a given point in space

■ Units: Hertz(Hz) or cycles(s-1) - (Hz=s-1) 

○ The speed of light is 2.9979 x 108 m/s 


What is the formula for plancks's constant?



○ The longer the wavelength, the lower the frequency, the lower the energy ○ The shorter the wavelength, the higher the frequency, the higher the energy ● Quantized Energy

○ Max Planck’s proposal (1918 Nobel prize): energy emitted or absorbed by an atom is quantized 

■ Atoms could absorb or emit energy only in certain discrete amounts ■ Quantum: smallest amount of energy

■ Energy of each quantum is proportional to its frequency

■ Energy formula: E=hv 

■ Planck’s Constant: h=6.626 x 10-34 

● Photoelectric Effect and Photons

○ Photoelectric Effect: electrons are emitted from a metal surface when light strikes it


What does emission spectrum indicate?



■ Light must have a certain minimum frequency (threshold frequency, vo) to eject electrons If you want to learn more check out What do puritans restrict other people to have?

■ When v<vo, no electrons are emitted, regardless of intensity of light ■ When v>vo, the amount of electrons increase with the intensity of light; and the kinetic energy of emitted electrons increases linearly with the frequency of light

■ When photon at threshold frequency: Eo=hvo 

■ When photon at a higher frequency: KEelectron=1/2mv2=hv-hvo 

○ Einstein’s Proposal: (1912 Nobel Prize) light acts like particles

■ Light consists of particles called photons

■ Energy of a photon is given by Planck’s relation: E=hv

● Atomic Spectrum of Hydrogen

○ White light gives a continuous spectrum when passed through a prism ○ Emission Spectrum: non continuous or line spectra

■ Indicates that the energy of the electron in the hydrogen atom is

quantized

● The Bohr Model: Bohr (1922 Nobel prize)

○ Electron in hydrogen atom moves around the nucleus only in certain allowed circular orbits 

○ Energy levels available to the electron in hydrogen atom

■ E= -2.178 x 10-18J x (Z2/n2) We also discuss several other topics like What is the most important aspect of crime control?

● n=integer

● Z=nuclear charge

○ Electrons emit radiation when they “jump” from an orbit with higher energy down to an orbit with lower energy

○ En= -(2.18 x 10-18J) 1/n2 

○ ��Eelectron= Efinal state- Einitial state 

○ ��E= -2.18 x 10-18J (1/nf2- 1/ni2) 

○ ��E=hv 

○ C=λv 

○ Pros:

■ Correctly fits the quantized energy levels of the hydrogen atom

● Suggests only certain allowed circular orbits for electron

■ As electron becomes more tightly bound, its energy becomes more negatively relative to the free electron

● The free electron is at infinite distance from the nucleus We also discuss several other topics like How does emile durkheim explain functionalism?

● As the electron is brought closer to the nucleus, energy is

released from system

○ Cons:

■ Only works for hydrogen

■ Electrons don’t move around nucleus in circular orbits

● Hydrogen Spectrum

○ Balmer series: nf=2, n i= 3,4,5,6 

○ Lyman series: nf =2, ni=2,3,4,5 If you want to learn more check out In what way are synapomorphies helpful to people?

○ Paschen series: n f=3, n i=4,5,6 

● Wave-Particle Duality of Electrons

○ De Broglie’s proposal (1929 nobel prize): particles could have wavelike character ○ Prediction: wavelength of a particle is inversely proportional to its momentum ○ λ=h/mv 

■ mv=momentum

■ For an electron 

● m=9.11x10-31 

● v=2.74x106 

● λ=2.64x10-10 

○ The wave and particle nature of the electron are complementary properties ■ The more you know about one, the less you know about the other ● Heisenberg’s Uncertainty Principle If you want to learn more check out How are simple monomers related to protobionts?

○ There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time

■ (��x) ��(mv)≥h/4π

● ��x=uncertainty in position

● ��v=uncertainty in velocity

● ��(mv)=uncertainty in momentum

■ The best we can do is to describe the probability that an electron will be found in a particular region using statistical functions We also discuss several other topics like What are the controversies about intelligence?

● Quantum Mechanical Model

○ Shrödinger’s Equation: Ĥψ=Eψ 

■ Ĥ=hamiltonian operator

■ E=energy of electron

■ ψ=wave function

■ This equation allows to calculate the probability of finding an electron with a particular amount of energy at a particular location in the atom 

■ Solutions to this equation produce many wave functions, ψ

■ A wave function is called an atomic orbital or orbital 

● Physical Meaning of Wave Function

○ The square of the function (ψ2), represented as a probability distribution, indicates the probability of finding an electron near a particular point in space ○ (a) Probability distribution for H1s orbital in three-dimensional space ○ (b) The probability of finding the electron at points along a line drawn from the nucleus outward in any direction

● Physical Meaning of Wave Function

○ A radial probability distribution graph plots the total probability of finding an electron in each spherical shell versus the distance from the nucleus

○ (a) Cross section of the hydrogen 1s orbital probability distribution divided into successive thin spherical shells

○ (b) Radial probability distribution plot

○ The size of the 1s orbital can be stated as the radius of the sphere that encloses 90% of the total electron probability

● Quantum Numbers: three parameters in the wave function for the hydrogen atom

numbers

Allowed values

Total number

Determines

orbital’s...

Principal quantum number, n

1,2,3,4...

Size and energy

Angular

momentum

quantum number, l

0 to (n-1)

n

shape

Magnetic

quantum number ml

-1, to +1

2l+1

orientation

○ When n = 1, l = 0;

■ n = 2, l = 0, 1;

■ n = 3, l = 0, 1, 2;

■ n = 4, l = 0, 1, 2, 3

Value of l

0

1

2

3

Letter used

s

p

d

f

● Orbital Shapes and Energies

○ The n determines size and energy of orbital 

■ The larger the value of n, the larger the orbital and higher the energy the orbital have

■ The l determines shape of orbital 

■ Nodes: areas of zero probability

○ Describing an Orbital

■ Each set of n,l, and ml describes one orbital

■ Orbitals with the same value of n are in the same principal level or principal shell

■ Orbitals with the same values of n and l are in the same sublevel or subshell

■ # of sublevels within a level = n 

■ # of orbitals within a sublevel = 2l + 1 

■ # of orbitals in a level = n2 

● Spin Quantum Number, ms 

○ Electron spin: a property of electron that rotates on its axis; there are two possible directions of rotation

○ The spin quantum number has only 2 allowed values: +½ and -½ ○ Pauli Exclusion Principle: no two electrons in the same atom can have the same four quantum numbers

○ Thus, a 1s orbital can “hold” two electrons:

■ Electron 1: n = 1, l = 0, ml = 0, and ms = +½

■ Electron 2: n = 1, l = 0, ml = 0, and ms = -½

○ An orbital can only hold two electrons, and they must have opposite spins because only two values of ms are allowed

● Four Quantum Numbers

● Polyelectronic Atoms: atoms with more than one electron

○ Electron correlation problem: since the electron pathways are unknown, electron repulsions cannot be accurately

○ Penetration effect: 

■ 3s electron penetrates to the nucleus more than 3p electron

● This causes the 3s electron to be attracted to the nucleus more

strongly than a 3p electron

○ Energy level in polyelectronic atom: Ens < Enp< End < Enf

● Energy splitting in Polyelectronic Atoms

○ For hydrogen atom or single electron systems:

■ Sublevels in each principal all have same energy

■ Orbitals with the same energy are said to be degenerate

○ For polyelectronic atoms

■ Energies of the sublevels are split

● 4s<4p<4d<4f 

● Caused by electron-electron interaction

● Early Periodic Table

○ Mendeleev’s periodic table: organized known elements of the time in a table format

● He arranged the rows so that elements with similar

properties would fall in the same vertical columns

● Contained some gaps, which allowed him to predict the

existence and properties of undiscovered elements

● Modern Periodic Table: lists elements in order of atomic number rather than mass

○ Periods: rows

○ Groups: columns

● Electron Configuration: description of the orbitals occupied by electrons ○ Orbital diagram: a square representing each orbital and a half-arrow representing each electron in the orbital

● Rules of Electron Configuration

○ Aufbaus’ principle: electrons ordinarily occupy orbitals of the lowest energy available

■ Energy order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc 

● Ex: Li has 3 electrons - 1s22s1 

● Ex: Ne has 10 electrons - 1s22s22p6 

■ When electrons occupy the lowest energy orbitals the atom is in its ground state

■ When electrons occupy any other orbitals the atom is in an excited state

○ Pauli exclusion principle: one atomic orbital can accommodate only 2 electrons, and these electrons must have opposite spins

○ Hund’s rule: electrons in the same subshell occupy the orbitals singly, with spins parallel, before they are paired

○ Condensed electron configuration:

■ (Z = 11) Na: 1s22s22p63s1 Na: [Ne]3s1 1s22s22p6=Ne

■ (Z = 3) Li: 1s22s1 Li: [He]2s1 1s2=H

● Valence Electrons and Core Electrons

○ For main group elements: the electrons in the highest principal energy level called valence electrons - blue

■ Valence electrons are the most important in chemical bonding ○ Electrons in lower energy shells are called core electrons - pink ■ Ex: Si 1s22s22p63s23p2 

■ Ge 1s22s22p63s23p64s23d104p2 

○ The elements in the same group on the periodic table have the same valence electron configuration and display similar chemical behavior ○ For main group elements - number of valence electrons=group number ● Orbital Blocks on Periodic Table

○ Main group elements:

■ S-block: ns subshell fills 

■ P-block: np subshell fills 

○ Transition elements:

■ D-block: (n-1)d subshell fills 

○ Inner transition elements: f-block: (n-2)f subshell fills 

● Period Trends in Atomic Properties

○ Atomic radius: one-half of the distance between covalently bonded nuclei

■ Increases in going down a group

● Orbital size increase in successive principal quantum levels ■ Decreases in going across a period from left to right

● Effective nuclear charge increases

● Valence electrons are drawn closer to the nucleus,

decreasing the size of atom

○ Ionization energy(IE): amount of energy required to remove an electron from the ground state of a gaseous atom or ion

■ Requires the input of energy to remove the electron

■ The lower the IE, the easier the electron is to be removed ■ It requires more energy to remove each successive electron ■ When all valence have been removed, IE takes a quantum “leap” ■ 1stionization: Mg (g) Mg+(g) + e-I1 = 735 kJ/mol

2ndionization: Mg+(g) Mg2+(g) + e-I2 = 1445 kJ/mol

3rdionization: Mg2+(g) Mg3+(g) + e-I3 = 7730 kJ/mol

■ In general:

● I1 decreases in going down a group

○ The electrons being removed are farther from the nucleus

● increases going left to right

○ Electrons added to the same principal quantum level

cannot completely shield the increasing nuclear charge

and are generally more strongly bound from left to right on

the periodic table

○ Electron affinity(EA): energy change associated with the addition of an electron to a gaseous atom

■ General trend for main group elements:

● Electron affinities become more negative from left to right

across a row

● The halogens have the most negative electron affinities

● Most groups of the periodic table do not exhibit any definite trend in electron affinity

○ Effective nuclear charge (Zeff) 

■ A net positive charge attracting a particular electron

■ Increases from left to right across a period

■ Doesn’t change much going down a group

● Metals, Nonmetals, and Metalloids

○ Metals: 

■ Low IE, small EA

■ Tend to form cations

○ Nonmetals:

■ High IE, large negative EA

■ Tend to form anions

○ Metalloids:

■ Elements that exhibit both metallic and nonmetallic properties ● Alkali metals 

○ Li, Na, K, Rb, Cs, and Fr have low ionization energy, and are the most chemically reactive of the metals

○ Hydrogen exhibits a nonmetallic character due to its small size

Page Expired
5off
It looks like your free minutes have expired! Lucky for you we have all the content you need, just sign up here