Ap Chemistry Unit 2: Atomic Structure

Wavelength

Wavelength

Crest to crest or trough to trough.

Frequency

The wave cycles that pass a given point.

Amplitude

Maximum height of a wave.

Node

Point of zero amplitude.

James Maxwell

All radiation propagates through space.

Radiation

λ x ν =c

Speed of Light

2.99 × 10^{8 m/s}

Diffraction

When waves have to bend around an obstacle.

Constructive Interference

When waves interact so that they add to make a larger wave, in phase.

Destructive Interference

When waves interact so that they cancel each other, out of phase.

Matter

A collection of particles.

Energy

A collection of waves.

Max Planck

Solved the “ultraviolet catastrophe.”

Photoelectric Effect

It was observed that many metals emit electrons when a light shines on their surface.

Classic Wave Theory

If the wavelength of light is made shorter or the light waves intensity made brighter, more electrons should be ejected.

Quantum Mechanics

An object can gain or lose energy by emitting or absorbing radiant energy not continuously but in discrete lumps or packs.

Photon

Streams of light particles (a quanta of light).

Quantum

Packet of energy.

Planck’s Constant, h

6.6262 × 10^-34J/s

Photon Formula

E=hv

Binding Energy

One photon at the threshold frequency gives the electron just enough energy for it to escape the atom.

Kinetic Energy Equation

E_k =1/2mv²

Bohr’s Major Idea

The level of energy in an atom is directly proportional to the distance from the nucleus.

De Broglie’s Proposition

If light can have material properties, then matter should have wave properties.

Principal Quantum Number

Tells you how far an electron is from the nucleus.

Orbital Quantum Number

Describes the shape of the orbital- s, p, d, f.

Orientation Quantum Number

S-1, p-3, d-5, f-7.

Electron Spin Quantum Number

The direction the electron is spinning.

Electron Configurations

Describes where all the electrons in an atom are located.

Aufbau Principle

Electrons will occupy the orbitals of lowest energy first.

Hund’s Rule

Electrons occupy the orbitals of the same energy with the same spin, until they have to double up.

When does promotion happen?

When the orbital is one away or half full.

Valence Electrons

Electrons in the highest principal energy shell.

Core Electrons

Electrons in the lowest energy shell.

Ferromagnetism

When something is always magnetic.

Paramagnetism

When something isn’t magnetic, but is with a magnet.

Diamagnetic

Something that’ll never be magnetic.

Atomic Radius

Half of the distance between the nuclei of 2 identical atoms joined in a molecule.

Coulomb’s Law

Explains that distance and attraction are directly proportional.

Group Trend

Atomic radii increases down the periodic table because of new energy levels.

Shielding Effect

Core electrons will shield valence electrons from the electrostatic force of the protons.

Periodic Trend

Atomic radii decreases across a period because more effective positive charge in the nuclei exerts a stronger pull.

Isoelectronic

Ions have the same number of electrons but their sizes vary.

Ionization Energy

Energy required to remove one electron from an atom that is a gas.

Electron Affinity

Energy change accompanying addition of electron to gaseous atom.

Bond Theory

Explains how and why atoms attach together, why some things combine, and stability.

Lewis Structure

Quick way to show a molecule’s shape, bond length, and polarity.

Exceptions To The Octet Rule

• Hydrogen- duet.

• Beryllium- quartet.

• Boron- sextet.

Resonance Structure

When there is more than one lewis structure for a molecule that differ only in the position of electrons.

Formal Charge

Fictitious charge assigned to help determine if the arrangement is likely (the best structure).

Formal charge Equation

Group # - (# of bonds+electrons).

Dipole

A polar bond.

Polar Covalent Bond

Charge distribution is NOT equable resulting in a dipole.

Polar Molecules

Must have dipole or polar bonds, and not be a “symmetric” shape.

Bond Order

The number of bonding electron pairs shared by two atoms.

Bond Order Equation

Number of shared pairs/atoms bonded to the central atom.

Bond Length

Distance between the nuclei of bonded atoms.

Isomers

The same atoms arranged differently can have different polarities.

Bond Energy

The energy change for breaking bonds in a molecule that is in the gaseous state.

Endothermic Reaction

A reaction that takes energy.

Exothermic Reaction

A reaction that releases energy.

Lattice Energy

Energy evolved when ions in the gas phase come together to form 1 mole of solid crystal.

Hess’s Law

If a reaction is the sum of two or more reactions, the overall change in energy is the sum of the change in energy for each reaction.

Valence Bond Theory

Atomic orbitals overlap and make new molecular orbitals.

Sigma Bond

Orbital overlap is between plane of nuclei.

Pi Bond

Orbital overlap is not in the plane of nuclei.

What Forms a Double Bond?

A sigma and a pi bond.

What Forms a Triple Bond?

A sigma and 2 pi bonds.

Organic Chemistry

Chemistry dealing with carbon compounds.

Hydrocarbons

Just hydrogen and carbon.

Aliphatic

Open chains.

Cyclic

Closed chains.

Saturated

Organic compound full of hydrogens (single bonds).

Unsaturated

Organic compound containing one or more multiple bonds.

Conjugated

Alternating single and double bonds.

Aromatic

Conjugated cycloalkenes.

Delocalization

Electrons move throughout the compounds.

Alkanes

Single bonds (-ane).

Alkenes

Double bonds (-ene).

Alkynes

Triple bonds (-yne).

Dienes

Two double bonds (-diene).

Cycloalkanes

Single bonds in a ring (cyclo-root-ane).

Cycloalkenes

Double bonds in a ring (cyclo-root-ene).

MO Theory

Overlap of atomic orbitals (wave functions) creates constructive (bonding) and destructive (antibonding) interference.

Electron Density

Where the electrons are likely to be.