Chem ch 6 test

Peaks

high points

Troughs

low points

Wavelength

distance between peaks/troughs (λ)

Amplitude

Half the distance between peak and trough

Frequency

the number of wavelengths that pass a given point in one second

Wavelength and frequency relationship

inversely proportional; as one increases the other decreases; If a wave has a very large wavelength, this necessarily means that the frequency will be short and vice versa.

Wavelength formula

λ = c/v (wavelength = speed of light/frequency)

Electromagnetic radiation

the means by which energy is transmitted through space

Speed of light

the speed of light, 'c', which is 3.00 x 10^8 m/s.

Blackbody radiation

refers to the radiation emitted by an object that is heated to extremely high temperatures.

Quantized

only certain allowable values for the energy

Photoelectric effect

Shining light of a specific frequency (Energy) onto metal caused the metal to emit an electron.

Photon

a packet of energy

Energy of a photon

E = hν or E = hc/λ

Energy formula

E = hv (E = planck's constant X frequency)

Line Spectra

a series of sharp lines in a spectrum that is characteristic of a chemical element.

Excited state

an electron in an atom w/ excess energy

Ground state

an electron in an atom in the lowest possible state

Quantized energy levels

Electron moves in a circular orbit.

Electron transitions

Electrons jump between levels by absorbing or emitting a photon of a particular wavelength.

De Broglie

theorized that if light can behave as a particle, then perhaps particles can behave as waves.

De Broglie wavelength formula

λ = h/mv

Uncertainty principle

we can never know the position and momentum of an electron exactly.

Schrodinger's equation

An equation to describe an electron's wave like properties.

Wave function (Ψ)

A mathematical function that describes the quantum state of a particle.

Nodes

Points where the amplitude of the wave has canceled out.

Antinode

Points where the amplitude of the wave is maximized.

Probability density/electron density

(Ψ2) gives the probability of an electron's location.

Atomic orbital

The area of space that corresponds to a 90% chance of finding an electron.

Principal quantum number (n)

Tells us how far away from the nucleus an electron is and is directly proportional to the energy of the electron.

Angular momentum quantum number (l)

Tells us the shape of the orbital in three-dimensional space.

Magnetic quantum number (ml)

Describes the orientation of the orbital in three-dimensional space.

s orbitals

Spherical in shape, with only one orbital.

p orbitals

Shaped like 'dumbbells' with three possible orbitals directed along each coordinate axis.

Nodal plane

A planar region where there is a zero probability of finding an electron.

d orbitals

Five orbitals that are important for explaining and predicting the reactivity of transition metal compounds.

Aufbau principle

In the ground state of an atom, electrons fill subshells of the lowest available energy first.

Electron Spin Quantum Number (Ms)

Has only two allowed values: + 1/2 and - 1/2.

Pauli Exclusion Principle

No two electrons in an atom can have the same four set of quantum numbers.

Hund's rule

The p electrons occupy the p orbitals singly, rather than pairing in one orbital.

Electron configuration

Shows the occupation of orbitals by electrons for a particular atom (e.g., 1s^1, 2s^2).

Orbital diagram

Electron configuration represented in boxes; two in the same box must have opposite spins.

Condensed electron configurations

Put the closest noble gas and then continue the configuration.

Valence electrons

Electrons in the outermost principal energy level involved in chemical bonding.

Core electrons

Electrons that are not in the outermost principal energy level.

Max electrons in S orbitals

2 electrons max (1 orbital per subshell).

Max electrons in P orbitals

6 electrons max (3 orbitals per subshell).

Max electrons in D orbitals

10 electrons max (5 orbitals per subshell).