General Chemistry

Matter

Anything that occupies space and has mass

Pure Substances

Substances that have fixed composition and properties

Pure Substance Examples

Water, iron, oxygen, sodium chloride

Elements

Pure substances that contain only one kind of atom

Element Examples

Magnesium, calcium, phosphorus

Compounds

Pure substances that contain two or more atoms that are chemically bonded together

Compound Examples

Water, sodium chloride, sucrose

Mixtures

Results when two or more pure substances are mixed

Mixture Examples

Sea water, sand, coffee

Homogenous Mixtures

Solutions; mixtures that show uniform composition where individual substances cannot be distinguished

Homogenous Mixture Examples

Sea water, sugar dissolved in water

Heterogenous Mixture

Mixtures where the composition varies and substances can be distinguished

Heterogenous Mixture Examples

Mixture of sand and seawater

Suspension

Heterogenous mixtures where one of the component particles is large enough to settle at the bottom

Suspension Examples

Mixture of sand and seawater

Colloid

Heterogenous mixtures where the components do not settle at the bottom

Colloid Examples

Mayonnaise, milk

Stoichiometry

Section of chemistry that deals with the quantitative relationship of the reactants and the product of a reaction

Theoretical Amount of Product

Can be predicted using stoichiometry given a particular amount of reactant

Amount of Reactant Needed

Can be determined using stoichiometry for a desired quantity of product

Moles

SI unit for amount of substance

Avogadro's Number

6.022 x 10^23 particles (equal to one mole)

Molar Mass

The ratio between the mass of substance for every mole of the substance

Molar Mass Source

Found for every element in the periodic table (atomic weight)

Molar Mass of Compounds

Sum of the atomic weights of the atoms present in the compound

Molar Mass Usual Unit

g/mole

Molar Mass of CO2

44 g/mole

1 mole of CO2 Mass

44 grams

1 mole of CO2 Molecules

6.022 x 10^23 molecules of CO2

Formula for Moles (n)

n = mass / MW

Mass Unit in Molar Formula

Grams (g)

MW Unit in Molar Formula

g/mol

Law of Conservation of Mass

Law stating that in a chemical reaction, matter is neither created nor destroyed

Reactant vs Product Mass

Total mass of reactants should equal total mass of products

Reactant vs Product Atoms

Number of atoms of an element in reactants equals number of atoms of same element in products

Requirement for Balancing Equations

Must be balanced in accordance with the Law of Conservation of Mass

Step 1 in Balancing Equations

Count the atoms of each element present on reactant and product sides

Step 2 in Balancing Equations

Adjust the coefficient of the substance accordingly if atom counts are not equal

Coefficient Definition

A whole number placed before a substance to indicate the number of moles present in the reaction

Step 3 in Balancing Equations

Check if the reaction is balanced; if not, repeat the process

Step 1 of Stoichiometric Relationship

Convert the given quantity to moles

Step 2 of Stoichiometric Relationship

Use the balanced equation to determine stoichiometric relationship between given and asked quantity

Step 3 of Stoichiometric Relationship

Determine the number of moles of the quantity being asked

Step 4 of Stoichiometric Relationship

Calculate the mass/volume of the quantity being asked

Solutions

Homogeneous mixtures composed of solute and solvent

Solute

The component in a solution that is less compared to the other

Solvent

The component in a solution that is greater compared to the other

Saturated Solutions

Solutions that hold the maximum amount of solute possible

Unsaturated Solutions

Solutions that contain solute whose amount is less than the maximum possible

Supersaturated Solutions

Solutions that contain more than the maximum amount of solute possible

Supersaturated Solution Stability

Unstable; excess solute precipitates or crystallizes when disturbed (agitation, adding crystals, scratching walls)

Like Dissolves Like Principle

Polar/ionic solvents dissolve polar/ionic solutes; nonpolar solvents dissolve nonpolar solutes

Exothermic Dissolution and Temperature

Increasing temperature decreases solubility

Endothermic Dissolution and Temperature

Increasing temperature increases solubility

Gas Solubility and Temperature

Increasing temperature decreases solubility

Solids/Liquids Solubility and Pressure

Not significantly affected by pressure

Gas Solubility and Pressure

Increasing pressure increases solubility

Concentration Definition

The ratio of solute and solvent in a solution

Percent Mass (%m/m) Formula

(mass solute / mass solution) x 100

Mass of Solution Formula

mass solute + mass solvent

Percent Volume (%v/v) Formula

(vol solute / vol solution) x 100

Volume of Solution Formula

vol solute + vol solvent

Percent Mass/Volume (%m/v) Formula

[mass solute (grams) / vol solution (mL)] x 100

Molarity (M) Formula

n solute / V (in L)

Molality (m) Formula

n solute / kg solvent

Mole Fraction (X) Formula

n solute / (n solute + n solvent)

Parts Per Million (ppm) Formula

(mass solute / mass solution) x 10^6

Parts Per Billion (ppb) Formula

(mass solute / mass solution) x 10^9

Normality (N) Formula

equivalents of solute / Vol solution (in L)

Equivalent of Solute Formula

mass (in g) / Equivalent Weight

Equivalent Weight Formula

Molar Mass / f

f Factor for Acids

Number of exchangeable H+

f Factor for Bases

Number of exchangeable OH-

f Factor for REDOX

Number of e- released or accepted

f Factor for Salts

Number of unique particles produced

Normality from Molarity Formula

N = M x n (where n is normality factor)

Dilution Definition

Addition of solvent in a solution to decrease concentration

Dilution Equation

C1V1 = C2V2

C1 and V1 Meaning

Initial concentration and original volume of solution

C2 and V2 Meaning

Final concentration and final volume of solution

Titration Definition

Analytical technique where a solution of known concentration determines an unknown concentration

Volumetric Analysis Titration Formula

nM1V1 = M2V2

n in Titration Formula

Stoichiometric factor

Colligative Properties Definition

Properties of solution that depend on the concentration of the solute

Boiling Point Elevation Definition

Addition of non-volatile solute increases boiling point compared to pure solvent

Boiling Point Elevation Formula

ΔTb = i x Kb x m

Tb (solution) Formula

Tb (pure solvent) + ΔTb

Freezing Point Depression Definition

Addition of non-volatile solute decreases freezing point compared to pure solvent

Freezing Point Depression Formula

ΔTf = i x Kf x m

Tf (solution) Formula

Tf (pure solvent) - ΔTf

Vapor Pressure Lowering Definition

Addition of non-volatile solute pulls vapor back to solution, decreasing vapor pressure

Vapor Pressure Lowering Formula

ΔP = Xsolute x P°solvent

Osmotic Pressure Definition

Minimum pressure required to oppose flow of solvent through semi-permeable membrane

Osmotic Pressure Formula

π = i x M x R x T

R Constant Value in Osmotic Pressure

0.0821 L atm/mol K

Temperature Unit in Osmotic Pressure

Absolute temperature in Kelvin

Van't Hoff Factor (i) Definition

Theoretical number of particles the solute produces after dissolution

Van't Hoff Factor of NaCl

2 (dissociates into Na+ and Cl-)

Kinetic Molecular Theory of Gases Definition

Theory explaining gas behavior based on particles in constant motion