Fundamentals of Chemistry and Principles of Stoichiometry (DONE)

SI system:

  1. Mass: kilogram, kg
  2. Temperature: kelvin, K
  3. Volume: litre, L
  4. Length: metre, m
  5. Time: second, s

Dimensional analysis (a.k.a. factor-label method):

  • Two physical quantities are equal if they have the
    same dimensions.
  • Technique for converting units of measurement into other units of measurement.
  • Uses conversion factors that are made from equalities between units.
  • Conversion units are arranged in fraction form in such a way as to cancel all other units except the desired unit.

Significant figures: Digits in a measured number that include all certain digits plus a final uncertain one.

  • Counting significant figures:
  1. All non-zero digits are significant
  2. Interior zeros are significant
  3. Trailing zeros after a decimal point are significant.
  4. Leading zeros are not significant
  5. Zeros at the end of a number without a written decimal point are not significant
  • Matter: Any object that has mass and
    volume
  • Mass: Amount of matter in an object
  • Weight: Force exerted by an object due to gravity (W = mg)

States of matter:

  1. Solid: Rigid matter, relatively incompressible, fixed shape and volume.
  2. Liquid: Fluid matter, relatively incompressible, fixed volume but not fixed shape (shape conforms to container).
  3. Gas: Fluid matter, easily compressible, no fixed shape (fills whole container).
  4. Plasma: Superheated ionized gas consisting of charged particles. No distinct shape.

Matter:

  1. Pure substances - constant composition, distinct property
  • Compounds - build up from basic particle of molecules of different atoms
    • Molecular compound
    • Formula units
  • Elements - one type of atom only
    • Atomic element
    • Molecular element
  1. Mixtures - combination of two or more substances which retain their distinct identities (can be separated by physical means)
  • Homogeneous mixtures - composition of the mixture is uniform throughout
  • Heterogeneous mixtures - composition is variable or not uniform (consists of regions called phases that differ in properties)

Physical changes: Changes that alter the state or appearance of the matter without altering the
composition.

  • Results in a different form of the same substance.

Chemical changes (a.k.a. chemical reaction): Changes that alter the composition of the matter.

  • The atoms that are present rearrange into
    new molecules, but all the original atoms are still present.
  • Results in one or more completely new substances. The new substances have different molecules than the original substances.
  • Different physical properties will be observed because the new substances have their own physical properties.
  • Separate mixtures based on different physical properties of the components:
  1. Different boiling point - Distillation
  2. Different state of matter - Filtration
  3. Different adherence to a surface - Chromatography
  4. Different volatility - Evaporation
  5. Different density - Centrifugation and decanting
  • Qualitative observations of reactions — changes in color and physical state.
  • Quantitative measurements — involve numbers (SI units are used).

Scientific notation: A way of expressing really big or really small numbers. (e.g. 10^12; 10^-9)

  • When multiplying or dividing measurements with significant figures, the result has the same number of significant figures as the measurement with the fewest number of significant figures.
  • When adding or subtracting measurements with significant figures, the result has the same number of decimal places as the measurement with the fewest number of decimal places.
  • Accuracy: The closeness of a measurement to
    its true value
  • Precision: The closeness of a set of values
    obtained from identical measurements of a
    quantity

Rounding off:

  1. If the digit is 5 or greater, round off to the last digit to be retained and delete all digits to the right.
  2. If the digit is less than 5, simply delete all digits to the right.

Components of atoms:

  1. Proton (p) - 1 a.m.u, +1 charge
  2. Electron (e) - 0 a.m.u, -1 charge
  3. Neutron (n) - 1 a.m.u, 0 charge
  • a.m.u = atomic mass unit

General symbol of an atom:
image

  • X: symbol for the atom
  • A: mass number (n + p)
  • Z atomic number (p)

Sizes of atoms:

  1. Volume
  • Assumed spherical and mentioned in atomic radius
  • Atomic volume increases as atomic number of atom increases
  1. Mass
  • 1 amu - 1.6605 x 10^-24 g
  • Molar mass of atom - mass of 1 mol of atom (g/mol)
  • Formula mass - Mass of 1 mol of molecule/formula (g/mol)
  1. Atomic mass: Mass of an atom measured in atomic mass unit (amu).
  • Exactly equal to one-twelfth the mass of carbon-12 atom.
  1. Atomic weight (a.k.a. average atomic mass): Average of the atomic masses of the different isotopes of an element, defined as the weight equivalent to 1/12 of the weight of one 12C atom (in amu or atomic mass unit)
  • Calculated as: Atomic wt = Σ(% Abundance x Mass of isotope)

Isotopes: Atoms that have the same number of protons and electrons but different numbers of neutrons and therefore have different physical properties (same atomic no., different mass no.)

Rules for naming simple chemical compounds:

  1. Names of compounds containing two elements end in –ide. The symbol of the cation (element that has undergone oxidation) is placed first.
  2. Acids that only contain two elements are usually named hydro....-ic acids.
  3. Acids with multiple elements have names that may have to be memorised.
  4. Radicals are groups of atoms that stay together in many chemical reactions, and act like single atoms. These have special names. (e.g. bicarbonate, hydroxyl, phosphate, carbonate, nitrate, sulfate)
  5. When two elements form a series of compounds, the prefixes mono-, di-, tri- etc. meaning 1, 2 or 3 atoms of the element in combination may be used to distinguish them.
  6. Many compounds have three elements (One of the three is usually oxygen, which is part of a radical) The ending –ate tells us that the compound contains more oxygen than a related compound that ends with –ite.
    • The prefix per- means more of an element than is found in another compound of those same elements
    • The prefix hypo- means less of an element than is found in another compound of those same elements
  7. For elements with more than one Oxidation State, Roman numerals are used to designate oxidation number.

Mole (also spelled mol): A standard scientific unit for measuring large quantities of very small entities such as atoms, molecules, or other specified particles.

  • 1 mole = the same number of atoms as there are in 12.00 grams of 12C
  • The no. of entities composing a mole has been experimentally determined to be 6.02214179 x 10^23, a fundamental constant named Avogadro’s no. or the Avogadro’s constant.
  • Counting: 1 mol = 6.022 x 10^23 units
  • Mass: 1 mol = mass formula of a compound in gram
  • Volume: 1 mol = 22.4L (any gas at 0°C and 1 atm)

Moles of constituent elements - examples:

  • 1 mol NaCl = 1 mol Na, 1 mol Cl
  • 1 mol H2O = 2 mol H, 1 mol O
  • 1 mol C6H12O6 = 6 mol C, 12 mol H, 6 mol O

Formula mass: The sum of the atomic masses of the atoms shown in the formula (generally used for ionic compounds).

  • Molecular mass is the same as formula mass but reserved for molecular substances.
    • Molecular mass expressed in units of gram per mole = molar mass.
  • The molecular formula is a multiple of the empirical formula.
  • To determine the molecular formula, the empirical formula and the molar mass of the compound has to be known.

Percent composition: Percentage of each element in a compound (by mass).

  • Can be determined from:
    1. The formula of the compound.
    2. The experimental mass analysis of the compound.
  • The percentages may not always total to 100% due to rounding.
  • Percentage = (part/whole) x 100%
  • Percentage = (mass of element X in 1 mol/mass of 1 mol of the compound) x 100%

The mass percent tells you the mass of a constituent element in 100g of the compound.

  • The fact that NaCl is 39% Na by mass means that 100g of NaCl contains 39g Na. (100g NaCl ≡ 39g Na)

Mole fraction (of component A, for example): Xa = Moles of component A/Total moles of all components

Stoichiometry: A study of the numerical relationship between chemical quantities (weight, mole, percent) in a chemical reaction.

  • Helps to predict how much of a reactant participates in a chemical reaction, how much product you'll get, and how much reactant might be left over.
  • The amount of every substance used and made in a chemical reaction is related to the amounts of all the other substances in the reaction (Law of Conservation of Mass).
  • Balancing equations by balancing atoms.
  • A chemical equation is an expression of a chemical process.

General information found in a chemical equation:

  1. Formula or symbols of reactants and products.
  2. States of the species involved in the reaction - (s) for solid, (l) for liquid, (g) for gas.
  3. Conditions for the reaction to occur (written above and below the arrows)
  4. Coefficients which indicate the relative proportions of all species in a reaction.
    • The coefficients can be used to relate all the species in terms of molecules and atoms, number of moles or weight by using the symbol '≡' which stands for equivalence.

Chemical equations must be balanced before using coefficients.

  • Balancing a chemical equation is essentially done by trial and error by counting the number of atoms in an equation.

Limiting reagent: The reactant which is used up before the other reactants in a reaction.

  • Is based on the stoichiometry of the reaction and the amount of reagent determines the amount of products obtained in the reaction.

Percentage yield = (Actual yield/Theoretical yield) x 100%

  • In order to determine the theoretical yield, the reaction stoichiometry should be used to determine the amount of product each of the reactants could make.
  • Theo theoretical yield will always be the least possible amount of product.
    • The theoretical yield will always come from the limiting factor.
  • Because of both controllable and uncontrollable factors, the actual yield of product will always be less than the theoretical yield.

Types of chemical formulae:

  1. Molecular formula: A formula which describes the exact composition of a molecule.
  2. Empirical formula: A formula which gives the smallest whole numbers that describes the ratios of atoms in a substance.
  3. Structural formula: A formula which describes the positions of atoms in a molecule.

Ionic compounds:

  • Made of ions called cations (+ charged ions) and anions (- charged ions).
  • The sum of the + charges of the cations must equal the sum of the - charges of the anions.
  • Writing formulas:
  1. Write the symbol for the metal cation and its charge.
  2. Write the symbol for the nonmetal anion and its charge.
  3. Charge (without sign) becomes the subscript for the other ion.
  4. Reduce subscripts to the smallest whole-number ratio.
  5. Check that the sum of the charges of the cation cancels the sum of the anions.
  • Most elements occur in nature as mixtures of isotopes.
  • The presence of all isotopes in an atom are in specific ratio known as abundance.
  • The atomic mass of atom is an average of the mass of all isotopes of the atom.
  • Ionic crystal: One giant crystal lattice.
  • Covalent molecules: Molecules with intermolecular force.

Combustion analysis: When a compound containing carbon and hydrogen is combusted, all carbon is converted to carbon dioxide and all hydrogen is converted to water.

  • The masses of carbon dioxide and water can be used to determine the amounts of C and H in the original compound.

The concentration of a solution is the ‘strength’ of a solution

  • The solute (a solid substance) is dissolved in the solvent (a liquid) to form the solution (the process is called solvation)
  • Most solutions are aqueous (aq) which means that the solvent is water. In some cases, the solute or solvent may also be a solid, liquid or gas.

A solution may be prepared by dissolving a more concentrated solution into the solvent, a process known as dilution.

  • Amount of solute stays the same but the concentration decreases.
  • Formula: M1V1 = M2V2
  • Apparatus used for dilution: Volumetric flasks, pipets.
  • Dilute solutions: Low amounts of solute per amount of solution.
  • Concentrated solutions: High amounts of solute per solution.

Density, ρ = Weight of solute/Volume of solution

  • Mass per unit volume of a substance

Weight percent (percent by weight): The mass of solute that makes up 100 units of mass.

  • % Weight = (Weight of solute/Weight of solution) x 100
    • OR: A (%w/w) = A g solute/100 g solution
  • Designated unit is (%w/w)

Volume percent (percent by volume): The volume of solute that makes up 100 units of volume of a solution.

  • % Volume = (Volume of solute/Volume of solution) x 100
    • OR: A (%v/v) = A mL solute/100 mL solution
  • Designated unit is (%v/v)

Parts per million (ppm): 1 ppm contains -

  1. 1 g of solute for each million grams of solution
  2. 1 mg of solute per kilogram of solution
  3. 1 mg of solute per litre of solution
  • Weight in volume

    ppm = grams of solute/10^6 milliliters of solution

  • Weight in weight

    ppm = grams of solute/10^6 grams of preparation

  • Volume in volume

    ppm = milliliters of solute/10^6 milliliters of solution

Parts per billion (ppb): The number of parts in one billion parts of the whole solution.

  • ppb = (Weight of solute/Weight of solution) x 10^9
    • OR: A ppb = A g solute/10^9g solution
  • 1 ppb
    • 1 g of solute per billion grams of solution
    • 1 μg of solute per liter of solution
  • Weight in volume

    ppb = grams of solute/10^9 milliliters of solution

  • Weight in weight

    ppb = grams of solute/10^9 grams of preparation

  • Volume in volume

    ppb = milliliters of solute/10^9 milliliters of solution

Molarity, M: Moles of solute per 1 liter of solution.

  • Used because it describes how many molecules of solute in each liter of solution.
  • Units: mol/L or molar or M
  • Molarity = (Moles of solute/Liters of solution)
    • OR: A molar = (A mol solute/1 liter solution)
  • When strong electrolytes dissolve, all the solute particles dissociate into ions.
  • By knowing the formula of the compound and the molarity of the solution, it is easy to determine the molarity of the dissociated ions. Simply multiply the salt concentration by the number of ions.

Molality, m: Number of moles of solute per kg of solvent.

  • Does not vary with temperature because based on masses NOT volumes.
  • Units: mol/kg or molal or m
  • Molality = (Moles of solute/kg of solvent)

Titration: A technique that uses reaction stoichiometry to determine the concentration of an unknown solution.

  • Titrant (unknown solution) is added from a buret.
  • Indicators: Chemicals that are added to help determine when a reaction is complete.
  • The endpoint of the titration occurs when the reaction is complete.

Acid-base titration:

  • The base solution is the titrant in the buret.
  • As the base is added to the acid, the H+ reacts with the OH- to form water. But there is still excess acid present so the colour of the indicator does not change.
  • At the titration's endpoint, just enough base has been added to neutralise all the acid. At this point, the indicator changes colour.

Mole equivalent relationship of the acid and the base: MaVa/A = MbVb/B

  • Ma and Mb for molarity
  • Va and Vb for volume
  • A and B for stoichiometric reaction cofficients
  • Burette readings are recorded up to two decimal places with the last digit being either 0 or 5 only.
    • Readings which differ by values greater than 0.20 mL may have to be discarded when calculating the average volume of titrant.
  • Pipets and volumetric flasks are recorded without any decimal places.