Features of Waves (3.1)

Electromagnetic radiation
- a wave that travels at the speed of light
- the electric field is the sine wave and is measured in voltage
- magnetic field is the cosine wave
quantized waves
- fixed ends of stationary wave restrict number and types of waveforms
- only discrete values of a continuous set are allowed
- node - region where amplitude is zero
  - no motion

Energy, wavelength, frequency of a photon (3.1)

wavelength goes in the opposite direction of frequency and energy
frequency is represented by v
wavelength is represented by lambda
the visible light spectrum is 400-800nm
- blue is 400nm and red is 800nm
https://www.cyberphysics.co.uk/graphics/diagrams/waves/EMSpectrumcolor.png
one photon is equal to one light unit
Energy is measured in joules

Photoelectric effect (3.1)

minimum threshold to eject electrons is related to the frequency of light
- classical physics suggested the amplitude of a wave influences the energy but experimentally, it was discovered that it is the frequency
there is a certain amount of energy required to hold the atom together
in order to eject an electron, you must overcome the threshold energy

Orbit energy of e-, delta E, Bohr Model (3.1, 3.2)

line spectra
- elements and molecules have emission spectra that correspond to a series of line
  - each line is a single wavelength
  - Rydberg equation (included in sheet)
    - n refers to which orbit an electron might be in
    - 1 over lambda
  - https://i.stack.imgur.com/oDIES.jpg
wave particle duality
- light behaves both wavelike and particle-like
  - wavelike property - interference patterns obtained when light passed through a double slit
  - particle like property - photoelectric effect
Bohr's atomic theory
- Bohr used the results of Planck and Einstein's experiments to develop a better atomic theory, making the following assumptions about an atom's electrons
  - electrons travel in circular orbit around the nucleus
  - energy proportional to distance
  - orbits are quantized
    - equation is included on the sheet
    - E sub n
    - Delta E
    - Z is the nuclear charge
  - energy is absorbed/emitted when orbits change
- Orbit transitions
  - an electron moves to a higher level (n) when a specific quantity of energy is absorbed
  - an electron moves to a lower level (n) when a specific quantity of energy is emitted
- the Bohr model fails with more than one electron, but was right about the following
  - e- attracted to nucleus by the opposing charges of the particles
  - e- energies are quantized
  - e- energies increase as the distance from the nucleus increases
  - discrete energies in spectra of elements result from quantized energy transitions of e- between levels

DeBroglie Wavelength (3.3)

wavelength is equal to h over linear momentum
- linear momentum is equal to mass multiplied by velocity
quantization of angular momentum, L, is equal to n multiplied by h
wavelike properties are only important for small particles (electrons)
- Electron location and momentum
  - Heisenberg Uncertainty Principle
    - We can only know either the position or momentum of a particle at any given instant
    - we can not know both in one instant
    - consequence of wave-particle duality
      - there are limitations governing the motions of quantum particles
schrodinger
- expanded on de Broglie's work and developed a wave equation that described and electron as a wave function, represented by psi
  - psi are not physical waves but complex probability amplitudes
  - psi squared is equal to the probability distribution (e- distribution)

Quantum Theory, Quantum Numbers (3.3)

a series of four quantum numbers describe the energies, locations and other properties of an electron in a given atom
- principle quantum number (n)
  - discrete energy level
  - energies increase as the distance from the nucleus increase
  - atomic orbital
  - which orbital
  - corresponds to the row on the periodic table
- angular momentum quantum number (l)
  - shape
  - integers from 0 to n-1
  - s, p, d, and f orbitals are 0, 1, 2, and 3, respectively
  - when l is not equal to 0, the orbitals have directionality about the z axis described by the magnetic quantum number
- magnetic quantum number (m sub l)
  - orientation
  - integers from -l to l
  - number of degenerate orbitals in subshell
- spin quantum number (m sub s)
  - electronic spin
  - overall one magnitude, two quantized states
  - +1/2, -1/2
  - spin states degenerate unless in external magnetic field

Orbitals and Wave functions (3.3)

pauli exclusion principle
- no two electrons can have the same two sets of electrons
- they have different quantum numbers in some way or another
radial nodes
- there are certain distances from the nucleus at which the probability density of finding an electron located at a particular orbital is zero
- the number of nodes is equal to the energy level minus l minus 1

Electron Configurations (3.4)

the total number of electrons in an atom populate these orbitals following a set of rules
the end result is an electron configuration or orbital diagram
hund's rule
- electrons spread out within a sub shell and have the same spin
- filled and half filled subshells are especially stable
aufbau principle
- build up from lowest energy to highest
valence electrons
- the electrons in the outermost principle quantum level of an atom (largest value of n)
- the 3de- are not counted as part of the valence shell
  - full d and f shells are not considered part of the valence shell because they are not involved in reactions if they are full
  - therefore, determination of the number of valence e- for any main group atom is simply the group number
  - HOWEVER, atoms in the transition series do count the de- toward the valence shell

Periodic Trends (3.5, 3.6)

periodic law - the properties of the elements are periodic functions of their atomic numbers
nuclear charge
- many of the periodic trends involving e- can be qualitatively explained using the effective nuclear charge (Z sub eff) of the atom
  - this is the pull exerted on a specific electron by the nucleus, taking into account any electron-electron repulsions
  - Z sub off is equal to Z minus the shielding
    - shielding is the e- repulsions
radii
- covalent radius - one half the distance between the nuclei of two identical atoms when they are joined by a covalent bond
- ionic radius - size of an ion
- increases as you go down and to the left
ionization energy - Energy required to remove the most loosely bound e- in formation of cation
- as size increases, IE decreases
  - it becomes easier to remove electrons
- increases as you go up and to the right
- watch for filled and half filled subshells
electron affinity - energy change for adding an electron to gaseous atom in formation of an anion
- increases going up and to the right
unit prefixes and conversions
- pico 10^-12
- nano 10^-9
- micro 10^-6
- milli 10^-3
- centi 10^-2
- kilo 10^3

Ionic and Covalent Bonds, including polyatomic (3.7)

monatomic ions
- the majority of chemical reactions do not change the nucleus of an atom
- electrons can be transferred or shared
- groups
  - alkali metals
  - alkaline earth metals
  - chalcogens
  - halogens
  - noble gases
polyatomic ions
- usually have a negative charge
binary bonding
- ionic
  - metal and nonmetal
  - groups of atoms held together by electrostatic interactions
  - transfer of electrons
- molecular/covalent
  - nonmetal and nonmetal
  - groups of atoms held together by sharing electrons

Nächster

Vorher

Chapter 1

Beschreibung