AP Chemistry - 01

Stoichiometry & Gases

The mole is a basic unit of measure in chemistry. One mole of a chemical has the same number of grams as the molecular weight of that substance.

For example, one mole of hydrogen (atomic weight = 1.008 amu) contains 1.008 grams of hydrogen.

One mole contains 6.02 x 10^23 individual units (atoms, molecules, etc.) of that substance. This is also called Avogadro’s Number.

The Mole

Chemical equations ( example: 2A + 1/2 B2 => C + D3 ) use the mole as the unit in the numerical coefficients and subscripts. Stoichiometric calculations ( ex: balancing equations ) will use and manipulate those numbers.

To write balanced equations, make sure that each element is present in the same number on each side, after taking into account subscripts and coefficients. Add coefficients, not subscripts, to balance.

Example: 2K3PO4 + 3Cs2Se => 3K2Se + 2Cs3PO4

The coefficient in front refers to the entire molecule. The subscript refers to that particular atom. The subscript with parentheses is for the marked group.

Naming and Identifying

Covalent compounds have numerical prefixes before the elemental names. ( Note that mono- does not attach to the first molecule)

Ex: H2O= Dihydrogen Monoxide, XeCl4 = Xenon Tetrafluoride

Ionic compounds are named [cation] [anion].

Ex: K+ / Cl- => KCl = Potassium Chloride

Individual ions are named according to the element, with certain suffixes, which depend on the # of O atoms.

NO- = nitride, NO2- = nitrite, NO3

- = nitrate

ClO- = hypochloride, ClO2- = chloride, ClO3

- = chlorate, ClO4

- = (hy)perchlorate

A hydrogen atom will raise the charge by one, and the ion is now [# prefix]hydrogen [original].

Ex: PO43- => H2PO4- : Phosphate => Dihydrogen phosphate

Isotopes / Mass Spectroscopy

Isotopes of an element have different numbers of neutrons, and thus different masses. Some are more prevalent than others. Isotopic abundance is used to calculate related data.

Mavg = x1M1 + x2M2, where x1 + x2 = 1, and x = % abundance of that isotope. Averaging them gives the average mass, which is given on the periodic table.

In a mass spectrometer (which is used to determine isotopic abundance, etc), atoms enter in a gaseous phase, and are accelerated through a uniform E-field, then deflected by a B-field. Particle mass determines the magnitude of deflection, and results are calculated using these data.

Gas Laws - Ideal

Formulae

PV = nRT, where:

P = Pressure (atm)

V = Volume (L)

n = # of moles of gas

R = .082 (L*atm)/(mol*K)

T = Temperature (K)

VRMS = (3RT/M)^1/2

V = root-mean-squared

R1/R2 = (M1/M2)^1/2

R = rate of effusion

The Kinetic-Molecular Theory assumptions are:

*Molecules are point masses

*Gas molecules exert no force on each other unless they collide

*Collisions are all perfectly elastic

*Gas molecules are in constant random motion

*The temperature of a gas depends on its average kinetic energy .

Note: diffusion = rate at which gas spreads. Effusion = rate at which gas escapes.

The most common velocity is the rms speed.

Gas Laws - Non-Ideal

The ideal gas equations start to break down when high pressure and low temperature cause intermolecular forces (they do exist)and molecular volume (it is not zero) to become issues, and violate the KMT.

* Pideal > Preal because intermolecular forces pull the molecules away from the walls.

* Videal < Vreal because molecular volume is non-zero and therefore takes up more space.

The revised equation is:

[ Preal + a(n/V)2)][Vreal -- bn] = nRT

[ R = .082, and a and b are determined experimentally ]

Pressure is the force that the gas molecules exert on the walls of their container, and is measured in mm Hg, Torr, (kilo)Pascal, atm, bar, and pound/square-inch. All are units of force per area.

In a mixture of gases, there is also partial pressure, which is the pressure that an individual gas in that mixture exerts on the container. Dalton’s Law states that the partial pressure of each gas is equal to (mole fraction)(total pressure). Total pressure equals the sum of all the partial pressures.

This is why one takes water vapor pressure into account when collecting data over water.

Pt = x1Pt + x2Pt + x3Pt + ...... + xnPt

Pt = P1 + P2 + P3 + ..... + Pn

Note that P1 = x1P1, and so forth.

Partial Pressure

1 atm = 760 Torr = 760 mm HG = 14.7 psi = 101.325 kPa

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