A. J.J. Thomson (1897) - Cathode ray experiment (discovered negative charges in an atom, electrons)

B. Rutherford Model - Alpha particle and gold foil experiment (discovered the dense center of an atom, nucleus)

A. One of the ways in which energy travels through space (anything with energy will emit electromagnetic radiation)

1. Wavelike behavior

a. Wavelength - the distance between two peaks (represented with lambda)

b. Frequency - number of waves that pass through a fixed point in one second (in hertz (1/second))

c. There is an inverse relationships between wavlength and frequency (long wavelength = low frequency)

i. Frequency and energy have a direct relationship (low frequency = low energy)

d. lambda x v = c (wavelength (m) times frequency (hertz) equals the speed of light (3.0 10^8 m/s))

B. In order of decreasing wavelength: radio waves, infrared, visible light, ultraviolet, x-rays and gamma rays

A. Max Planck (1858-1947) - Proved that absorbed energy was not released all at once but rather through packets over time

1. Energy can be gained or lost only in whole number multiples

2. Change in energy - delta-E = h x v (delta energy in joules (J) equals Planck's constant (6.626 10^-34 Jxs) times frequency (Hz))

3. Electromagnetic radiation can be viewed as a stream of particles called photons

a. Light can behave as a stream of photons or waves (dual nature of light)

b. Energy of a photon: E = hv (energy (J) equals Planck's constant (6.626 10^-34 Jxs) times frequency (Hz))

B. Photoelectric effect

1. No electrons are emitted by a given metal below a threshold frequency

2. For light with frequency lower than the threshold frequency, no electrons are emitted regardless of the intensity of the light

3. For light with frequency greater than the threshold frequency, the number of electrons emitted increases with the intensity of the light

4. For light with frequency greater than the threshold frequency, the kinetic energy of the emitted electrons increases linearly with the frequency of the light

5. Kinetic energy of an electron

a. KE electron = 1/2 m(v^2) (m is mass in kg)

i. 1/2 m(v^2) = hv - h(v sub 0)

ii. hv - energy of incident photon

iii. h(v sub 0) - energy required to remove electron from metal's surface

b. de Broglie's equation

i. lambda (m) = h (Planck's constant) / mv (mass in kg times speed in m/s)

A. As electrons absorb energy they jump up principal energy levels and return to their original principal energy level upon releasing the energy in packets

A. Dictates that electrons orbit around the nucelus in set levels

1. When the electron absorbs energy, the electron moves to the outer principal energy levels

2. It then travels back down the energy levels upon releasing the energy as a photon of light

B. Bohr Constant: -2.178 10^-18 J

1. E (J) = Bohr Constant (J) x (z^2/n^2)

a. Z is the number of protons and N is the principal energy level

2. If E is positive, it represents repulsion between electrons; if E is negative, it represents attraction to the nucleus from the electron

C. delta-E = E(final) - E (initial)

1. If delta-E is positive, energy is absorbed; if it's negative, energy is emitted

D. de Broglie's equation - lambda (m) = h / mv (kg, m/s)

A. There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time

1. delta-X x delta-MV > (or equal to) h/4pi

a. X is the particle's position, MV is the particle's momentum and h is Planck's constant (6.626 10^-34)

A. 7 principal energy levels (7 periods on the periodic table)

1. Maximum electrons for each level = 2n^2

B. Cu, Ag and Au borrow an electron from the previous S orbital to end in a D^10 orbital

1. Pd and Pt are also exceptions: they will borrow 2 electrons from the previous S orbital to end in a D^10 orbital

2. Cr and Mo are also exceptions: they will borrow 1 electron from the previous S orbital to end in a D^5 orbital

A. Atomic orbital - region of space around the nucleus of an atom where an electron is likely to be found

B. Set of quantum numbers

1. n (principal quantum number)

a. Describes the energy level on which the orbital resides (e.g. n = 1)

2. l (angular momentum quantum number)

a. Describes the shape of an orbital; l = only values from 0 to n-1 (e.g. n = 2, l = 0, 1)

b. l= 0 (s), 1 (p), 2 (d) and 3 (f)

3. m sub l (magnetic quantum number)

a. Describes the 3D orientations of the orbital; only values from -l to +l (e.g. n=2, l=0,1 then for l=0, m sub l=0 and for l=1 m sub l=-1, 0 and 1)

4. m sub s (spin quantumn number)

a. Describes the electron spin; an orbital can hold up to two electrons with opposite spin directions (e.g. -1/2 or 1/2)

i. Unreliant on other quantum number values, always written as +/- 1/2

C. Radial node - a sphere in an orbital with no electron probability

D. Radial probability - a graph for an orbital plotting the probability of finding the electron in it versus distance from the nucleus

1. To graph this, the x-axis is distance from the nucleus and the y-axis is the probability of finding an electron

a. Each energy level ascending has a higher chance of containing an electron than the previous level

i. As such, each peak on the graph is higher than the last

A. No two electrons in an atom can have the same quantum numbers

A. Electron shielding - when electrons are placed in a particular quantum level, they "prefer" the orbitals in the order s, p, d and then f

A. The energy required to remove an electron from a gaseous atom or ion where the atom or ion is assumed to be in its ground state.

1. e.g. Al(g) --> Al^+(g) + e^- in which I sub 1 = 580 kJ/mol

B. Electron affinity - energy change associated with the addition of an electron to a gaseous atom

1. X(gas) + e^- --> X^-(gas)

C. Atomic radii - defined as half the distance between the nuclei in a molecule consisting of identical atoms

D. Isoelectronic ions - ions with same number of electrons

1. e.g. S^2- has 16p and 18e^-; Ca^2+ has 20p and 18e^-

2. The more protons an isoelectronic atom has, the more effectively it can attract its electrons thus having a smaller radius

a. e.g. in the above example, Calcium 2+ has a smaller atomic radius than Sulfur 2-

E. Ionization energy increases from the bottom to the top of the periodic table, it also increases from left to right

F. Electron affinity also increases from bottom to top and left to right; however, chlorine has the highest electron affinity

G. Atomic radius increases from top to bottom and right to left

1. The more electrons an ion has, the larger the radius is

2. In two ions with different elements, the one with more protons has the smaller radius