Fundamentals of Chemistry

Fundamentals of Chemistry

• Chemistry studies the composition and transformations of matter.
• Matter is anything that has mass and occupies space.

Characteristics of Matter

• Possesses mass
• Is impenetrable
• Occupies space
• Is indivisible
• Cannot be destroyed

Law of Conservation of Matter

• Matter is neither created nor destroyed, it only changes.
• The mass before a chemical reaction equals the mass after the reaction.
• Remains constant, even with physical or chemical changes.

Energy

• Energy is the capacity to perform work or transfer heat.

Law of Conservation of Energy

• Energy is neither created nor destroyed, it only changes from one form to another.

Classification of Matter

• A pure substance is a simple or compound substance formed by identical atoms, molecules, or ionic agglomerates.

Simple Substance

• When composed of the same type of elements.

Simple Molecules

• H2, N2, O2, O3

Allotropes

• Are the same element, but appear in different forms.
• C of diamond - C of graphite

Compound Substance

• Consists of two or more different elements.
• Example: H2O

Mixtures

• Mixtures are combinations of two or more pure substances.
• Simple - Simple mixtures consist of two simple substances
• Simple – Compound mixtures consist of simple and compound substance

Homogeneous Mixtures

• Presents only one phase.

Heterogeneous Mixtures

• Presents two or more phases.

Methods for Separating Heterogeneous Mixtures

• Sieving (S - S): Separates by difference in size.
• Filtration (S – L or S - G): Separates solids from liquids or solids from gases.
• Decantation (S – L or L - L): For immiscible substances.
• Dissolution (S - S): Separates substances with different solubilities.
• Flotation (S - S)(resting): Separates by differences in densities.
• Levigation (S - S)(movement): Separates by differences in densities.
• Centrifugation (S – L or L - L): Separates substances of differing densities.
• Sublimation (S - S): One substance must be able to sublime to separate.
• Magnetization (S - S): One substance is attracted to a magnet.

Methods for Separating Homogeneous Mixtures

• Distillation: Simple (<80°C), Fractional (>80°C).
• Crystallization (S. dissolved - L).
• Electrophoresis (S - S): Separates substances with different charges.
• Chromatography (S – L o L - L): Stationary and mobile phases involved.

Properties of Matter

• Physical Properties: Changes the physical state, but not composition.
  • States of aggregation
  • Density
  • Viscosity
• Chemical Properties: Alters the chemical composition of the matter.
  • Oxidation
  • Reduction
  • Hydrolysis
  • Saponification
  • Digestion
  • Corrosion
  • Explosion
  • Inflammation
  • Combustion
  • Neutralization
• Intensive Properties: Non-additive; do not depend on the amount of matter.
  • Color
  • Odor
  • Taste
  • Boiling point
  • Melting point
  • Density
• Extensive Properties: Additive; depend on the amount of matter considered.
  • Weight
  • Mass
  • Volume
  • Concentration
  • Heat required

States of Aggregation of Matter

• Solid: Maintains constant shape and volume.
• Liquid: Variable shape, constant volume and adapts to the container's shape.
• Gaseous: Variable shape and volume; compressible and expandable.

Atoms, Molecules, and Ions

• Key experimental evidence for the existence of subatomic particles.

Electron: Cathode Ray Tube Experiments

• Cathode ray tubes, consisting of a vacuum tube with two metal plates connected to a high voltage source, emit an invisible ray from the negative plate (cathode).
• The cathode ray is directed towards the positive plate (anode), passes through a perforation, and continues to the end of the glass tube, producing fluorescence or bright light.
• Cathode rays are attracted to positive charges and repelled by negatives, indicating the existence of negatively charged particles now known as electrons.
• Thomson determined the ratio between electric charge and mass of an electron by using cathode ray tubes, it was found to be -1.76 × 10^8 C/g

Electron Charge: Oil Drop Experiment

• R. A. Millikan precisely measured the charge of an electron through experiments conducted between 1908 and 1917.
• Millikan analyzed the movement of tiny oil droplets that acquired static charge from air ions.
• The charge was determined to be 1.6022 × 10^-19 C.

Proton and Nucleus: Gold Foil Experiment

• In 1910, Ernest Rutherford used alpha particles to investigate the structure of atoms.
• Most particles passed through gold foil with little to no deflection.
• Some alpha particles were dispersed at large angles; some even bounced back.
• Rutherford proposed most of the atom is empty space, with positive charges concentrated in a central nucleus.
• The particles in the nucleus with a positive charge are protons.
• Protons have the same amount of charge as electrons, its mass is 1.67262 × 10^-24 g.

Neutron: Beryllium Plate Experiment

• Chadwick bombarded beryllium with alpha particles, it emitted a high-energy radiation similar to gamma rays.
• These rays consisted of neutral subatomic particles called neutrons with slightly more mass than the protons.

Atomic Number: Hydrogen Absorption Spectrum Experiment

• Moseley measured the frequencies of X-rays emitted by different elements.
• There is mathematical relation between an element's X-ray frequency and its atomic number.
• The periodic table should be ordered by atomic number and confirmed Bohr's model for X-ray energies and predicted undiscovered elements.
• The atomic number (Z) is the number of protons in the nucleus of an element.
• In a neutral atom, the number of protons equals the number of electrons, indicating the number of electrons in the atom

Mass Number

• The mass number (A) signifies the total number of neutrons and protons in an atom's nucleus.:
  • A = Z + N (A = mass number, Z = atomic number, N = number of neutrons)
  • Z = A - N
  • N = A - Z

Isotopes

• Atoms with the same number of protons but differing in the number of neutrons.
• Z = Z, same atomic number but different mass number.

Isobars

• Atoms of different elements with the same mass number (A=A).

Isotones

• Atoms of different elements with the same number of neutrons (N=N).

Atoms

• Defined as the basic unit of an element, maintains chemical identity through chemical/physical changes.
• Dalton described atoms as indivisible and extremely small.

Molecules

• Smallest particle of an element/compound that has independent and stable existence.

Diatomic Molecules

• Contain two atoms.
• N2, O2, F2, HCL

Polyatomic Molecules

• Contain more than 2 atoms.
• O3, H2O, NH3

Ions

• An ion is an atom or group of atoms with an electrical charge.
• Loss of electrons from neutral atom creates a positive charge.
• Gain of electrons creates a negative charge.
• Sodium chloride (NaCl), is an ionic compound formed by Na+ and Cl- ions and is known as table salt.
• Monoatomic ions contain only one atom.
• Polyatomic ions consist of multiple atoms.

Atomic Theory

• Democritus proposed that all matter is formed of small, indivisible particles names atoms.

Dalton's Atomic Theory

• Elements are composed of extremely small particles that are atoms.
• All atoms of the same element are identical in size, mass, and chemical property.
• Compounds are composed of atoms of more than one element.
• Chemical reactions involve separation, combination, or rearrangement of atoms, not their creation or

Periodic Table

• Elements organized by physical properties and behavior

• Elements are arranged by atomic number in horizontal rows, and by similar chemical properties in vertical columns(groups/families)

• elements are divided into metals, non-metals, and metalloids

• Metals are good conductors of heat and electricity

• Nonmetals are poor conductors of heat and electricity

• Metalloids have intermediate properties

• Group 1A (Li, Na, K, Rb, Cs y Fr) are alkali metals

• Group 2A (Be, Mg, Ca, Sr, Ba y Ra) are alkaline earth metals

• Group 7A (F, Cl, Br, I y At) are halogens

• Group 8A (He, Ne, Ar, Kr, Xe y Rn) are noble/rare gasses

Chemical Element

• Cannot be broken down by chemical reactions

Chemical Symbol

• Graphic representation of an element

Nomenclature for Binary Compounds

• Metal + Oxygen forms Basic Oxides
• Fe2O3 ((Sk) Iron(III) oxide, (T) Ferric oxide)
• Metal + O2 (simplified) forms Peroxides
• H2O2 (Hydrogen peroxide)
• Metal + Hydrogen forms Metallic Hydrides
• FeH2 ((Sk) Iron(II) hydride, (T) Ferrous hydride)
• Metal + Nonmetal forms Binary Salts
• FeCl2 ((Sk) Iron(II)Chloride, (T) Ferrous Chloride)
• Nonmetal + Oxygen forms Acid Oxides/Anhydrides
• I2O5 ((Sk) Iodine(V) oxide, (T) Iodic anhydride)
• Metal + (O2)1 forms Superoxides
• Sn1O8 ((Sk) Tin(IV) Superoxide, (T) Stannic Superoxide)
• Nonmetal + Hydrogen forms Hydrodes non metalics
• HCl ((Sk) Hydrogen chloride, (T) Hypochlorhydric acid)
• Nonmetal + Nonmetal forms Covalent Salts
• PCl3 ((Sk) Phosphorous(III) Chloride, (S) Phosphorus trichloride)

Ternary Compounds

• Metal + (OH) forms Hydroxides
• Al(OH)3 ((Sk) Aluminium(III) hydroxide, (T) Aluminum hydroxide)
• Anhydride + H2O forms Oxoacids
• H2SO4 (Sulfuric acid)
• (NH4) + Nonmetal forms Ammonium Salts
• (NH4)2S (Ammonium sulfide)
• Metal + (CN) forms Cyanide Salts
• Fe(CN)2 (Iron(II) Cyanide, (T) Ferrous cyanide)
• Metal + Nonmetal Oxide forms Ternary Salts
• Na(SO4) (Sodium sulfate)

Quaternary Compounds

• Acid Salts MHNmO: Na(HSO4) (Sodium acid sulfate)
• Basic Salts M(OH)Nm0:[Al(OH)]3 (PO4)2 (Monobasic aluminum phosphate)
• Double Salts MM NmO: NaAl(SO4) (Double aluminum and sodium sulfate)
• Ammonium: (NH4)NmO:(NH4)3 PO4 (Ammonium phosphate)
• Carbonates: (NH4)2 (CO3) (Ammonium carbonate)

Chemical Quantities:

• Uma: Atomic mass unit.
• Defined as the amount of substance or particles.
• Atomic mass: the mass that each of the isotopes present in nature has and is expressed in uma.
• Atomic weight: is the ponderal average sum of the atomic masses of the isotopes present in nature and is expressed in uma.
• Formula weight: is the sum of the atomic weights that constitute the formulas multiplied by the quantity that is repeated in uma. Mol: is defined as the amount of entities, such as the number of carbon atoms present in 12 g.