SCH4U - Unit Four Structures & Properties

JJ Thomson

used the cathode ray tube to discover electrons

Ernest Rutherford

used the gold foil experiment to discover the positively charged nucleus

Who developed the photoelectric effect?

Heinrich Hertz

photoelectric effect

objects will release e if struck by light

minimum threshold (energy) is needed

frequency (colour) determines the energy of the emitted e

Heinrich Hertz

the strength of emitted electrons depend on the frequency/colour of the light, not the brightness

Who developed quantum theory?

Max Planck

what did Max Planck study to develop quantum theory?

blackbody radiation

blackbody radiation

Energy (light/heat) given off by an object because of its temperature.

Hotter objects emit more energy and at shorter wavelengths (bluer)

quantization of energy

energy is not continuous, packaged in small amounts of quanta

Albert Einstein

light is both a wave AND particle

Albert Einstein's proposition

electrons were emitted from the surface of the metal because a photon collided with the electron

Why would an electron emit when colliding?

transferred energy to the electron would cause it to break away

Niels Bohr

determined electrons have certain energies

Bohr's discovery

the energy of electrons are quantized, existing in special energy states, ground and transition

transition/excited state

an electron can absorb energy and jump to a higher energy level (shell)

electron returns to its ground state

excess energy is emitted as photons

excess of energy emitted

different elements will produce different frequencies/colours of electromagnetic radiation

de Broglie

matter, specifically electrons, has particle-wave duality

Schrodinger studied

focused on the wave properties of electrons

Werner Heisenberg

uncertainty principle

uncertainty principle

wave and particle nature of electrons are complementary, being inversely related, the more we know about one, the less we know

uncertainty principle and electrons

since we cannot know both of these at a specific time, we state the probability of finding electrons

quantum model

electrons can be in different orbitals by absorbing or emitting energy, and the location of electrons is given by probability

orbitals

space where only two electrons can be found, based on probabilities

types of orbitals

S, P, D, F

S orbitals

spherical-shaped orbitals, n=1/L=0

P orbitals

Px, Py, Pz

number of D orbitals

five orbitals

number of F orbitals

seven orbitals

pauli exclusion principle

no two electrons can have the same set of quantum numbers

principal quantum number

describes size and energy level or orbitals, whole number value n = 1, 2, 3...

orbital shapes affected by principal quantum numbers

since they can exist at different energy levels, they increase in size with their energy

secondary quantum numbers

refers to the subshell and shape of orbitals, with whole number values from zero to n - 1

l = 0 (s)

l = 1 (p)

aufbau principle

lowest energy orbitals are filled first

Hund's rule

one electron is added to each level before electrons can be paired

electron configurations

summary of energy level diagram

principal number, orbital subshell, number of electrons

condense full electron configurations

by using noble gas placeholders

e.g, [Xe] ...

electron anormalies

half-filled subshells are more stabe than partially filled subshells

electron configurations differs

electrons can jump energy levels to go from partially filled, to half-filled subshell to become stable

charge affect configuration

add electrons to the lowest energy orbital (anion)

remove electrons from the highest principal number (cation)

isoelectronic

having the same electron configuration between to elements

paramagnetism

the weak attraction of a substance to a magnet, caused by unpaired electrons having the same spin

VSPER theory

method to determine geometry of molecules, based on how electron pairs will want to repel one another

electron groups on a central atom

groups of electrons will either be a lone pair, or other atoms connected to a central atom

triple bonds affect how many groups are considered

No...

single, double, and triple bonds are all considered as one group

AZE method

determine # of central atom: A

determine # of bonded atoms: Z

determine # of lone pairs: E

lone pairs

lone pairs take up more space, lessening the angles between atoms

Do you show angles between lone pairs and atoms?

No...

technically the lone pairs are invisible, only angles between atoms are considered