Atomic Structure

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Quantum Theory

Development of the Atomic Theories of Matter

Democritus - first had the theory, Greek philosopher

Dalton - revived the idea

J.J. Thompson - found positive and negative areas of the atom, using a cathode ray

Rutherford - found the nucleus, using alpha particles and gold foil (some bounced back)

Electromagnetic Spectrum

light behaves as a wave (moving through space) and a particle (interacting with matter)

as frequency increases, wavelength decreases

v = 3.0 x 10^8 m/s

v = λf

Planck's Theory

explaining line spectrums that come from heating gases

energy is absorbed/emitted in quanta (fixed amounts)

each colour of light corresponds to a quantum

higher frequency light has more energy than lower frequency light

Photoelectric Effect

when light shines on metal, electrons are ejected

requites a minimum frequency of light

light energy in photons - carries a specific amount of energy determined by frequency

photons transfer energy to electrons on surface - either all or none, one photon per electron

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The Dual Nature of Atoms

Bohr Model of Hydrogen

electrons "jump" up the quanta

when they fall down, they release quantized energy as photons of that particular wavelength (light)

problems - didn't account for repulsion between electrons

Matter Waves

matter has a dual nature, like light - exhibits wavelength

Heisenberg's Uncertainty Principle

position and momentum cannot be simultaneously measured and known exactly

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Quantum Model of the Atom

Quantum Model

models probability of location/speed of electron

The energy of electrons is quantized

Electrons exhibit wavelike behaviour

Cannot know exact position and momentum at the same time

Orbitals

no precise orbits - probability of finding an electron with a given energy level in an area around the nuclues

orbit - actual path electron follows

see charts/diagrams

Electron Spin

either clockwise or counterclockwise

creates magnetic field

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Electron Arrangement in Atoms

electron configuration - arrangement of electrons among the orbitals of an atom

The Aufbau Principle - electrons are added one at a time to lowest energy orbitals available

The Pauli Exclusion Principle - orbitals can hold a maximum of two electrons, must spin in opposite directions

Hund's Rule - only pair electrons once all the spots in the orbital have been filled

Sublevel Notation - 1s2 2s2 2p2, etc.

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Periodic Table Revisited

Basic Organization

atoms arranged in order of increasing atomic number

similar properties in vertical columns (groups); horizontal rows (periods)

Metals

all solid at room temperature, except mercury

shiny, malleable, ductile

good conductors

Non-Metals

many gases at room temperature, some solids, bromine is liquid

dull, brittle, hard or soft

poor conductors

Metalloids - some properties of each, or between

Valence Electrons

electrons that occupy the highest principal energy level are the valence electrons

elements in the same group have similar properties because the have valence electrons in similar configurations

Noble Gas Notation

atom's inner electrons are represented by the symbol for the nearest (lower) noble gas

eg. Lithium - [He]1s2

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Trends in the Periodic Table

Atomic Radius

distance from centre of radius to outermost electron

increases down - because principal energy level of outermost electrons increases - orbitals getting farther away

increases left - because amount of protons decrease, meaning weaker pull - larger radius

Ionic Size

losing electrons (positively charged) - atom becomes smaller

because - loss of valence electrons means outer orbitals are empty; loss of electrons means that there is more of a positive pull towards nucleus, less repulsive force between electrons

gaining electrons (negatively charged) - atom becomes bigger

because increase of repulsive force pushes electrons further apart

Ionization Energy

energy needed to remove one of its electrons

increases up and right, because further away from the nucleus means a weaker pull on the electrons being removed, and vice versa

every time an electron is removed, the next one is much harder

Electronegativity

how attractive an atom is to electrons

in general, increases up and right

non-polar covalent bond - sharing electrons basically equally, <0.4

polar covalent bond - one atom has a significantly higher EN than the other, causes partial chargers (delta +/-), 0.4 - 2.0

ionic - gives/takes, >2.0