General Chemistry Lecture Notes

I. INTRODUCTION TO MATTER

• GENERAL CHEMISTRY
  • Deals with the study of matter in terms of:
    • Composition
    • Structure
    • Properties
    • Changes
    • Energy
• Matter: Anything that occupies space and has mass.
  • Mass: Amount of matter present in an object
  • Volume: Amount of space occupied by an object
  • Weight: mass x gravity

PROPERTIES OF MATTER

• INTENSIVE/INTRINSIC PROPERTY
  • Properties that are independent of the amount of matter present
  • Examples:
    • Density
    • Specific gravity
    • Boiling point
    • Melting point
    • Organoleptic properties
• EXTENSIVE/EXTRINSIC PROPERTY
  • Properties that are dependent on the amount of matter present.
  • Examples:
    • Mass
    • Volume
    • Heat content
    • Pressure

CHANGES OF MATTER

• PHYSICAL CHANGE
  • Observed or measured without changing the identity of matter
  • Change in phase of matter
  • Examples:
    • MP, BP, Solubility, Miscibility, Viscosity
    • H2O(l) \rightleftharpoons H2O(s)
• CHEMICAL CHANGE
  • Describes the change or reaction a substance undergoes
  • Change in both intrinsic and extrinsic properties
  • Examples:
    • Flammability
    • Formation of Gas
    • Formation of Precipitate
    • Production of Odor
    • Change in color
    • Burning of paper

STATES OF MATTER

• STATES OF MATTER

• PLASMA
  • 4th state of matter
  • Also known as mesophase/liquid crystals
  • Properties:
    • Solid and flow-like properties
    • Crystal lattice – ordered arrangement of atoms
  • Main types:
    • Smectic – soap-like/grease-like
    • Nematic – thread-like
  • Crystal System

LAWS OF MATTER

• LAW OF DEFINITE PROPORTION
  • Also known as Proust’s Law/ Law of constant composition
  • States that the composition of a pure compound is always the same regardless of its source.
  • In pure compounds, the elements are always in the same proportion by mass
  • The same ratio must be a whole number
• LAW OF MULTIPLE PROPORTION
  • Also known as Dalton’s Law
  • Atoms of two or more elements may combine in different ratios to produce more than 1 compound.
  • The multiple ratios must be a whole number.
• LAW OF CONSERVATION OF MASS
  • No change is observed in the total mass of the substances involved in a chemical reaction

II. CLASSIFICATION OF MATTER AND ITS PROPERTIES

PURE SUBSTANCES

• Elements:
  • Simplest form of matter that has definite chemical compositions
  • Cannot be decomposed by simple physical/chemical means into two or more different substances

• Compounds:
  • Composed of two or more elements that unite chemically in different proportions.
  • Cannot be changed into simpler substances under normal conditions

MIXTURES

• Substances that are not chemically combined
• May be classified according to its:
  • Nature of particles: Homogenous vs Heterogenous
  • Particles size: Solution, Colloids, Suspensions
• NATURE OF PARTICLES

PARTICLE SIZES

• TRUE SOLUTIONS
  • Uniform mixture
  • Composed of solute + solvent
  • Atom, molecules, ions of the substance become dispersed
• COLLOIDS
  • Contains particles bigger than those in solutions but smaller than those in suspensions
  • Particles of solute are broken down to the size of molecules but are small enough and dispersed throughout the medium
  • Property of colloids:
    • Tyndall effect – light scattering effect
    • Brownian Movement – zig-zag movement of particles
    • Adsorption – capacity to adhere (stick) to the surface
    • Charged electrically – zeta potential
• SUSPENSIONS
  • Coarse mixture
  • Finely divided solid materials distributed in liquid

III. SOLUTIONS

FACTORS AFFECTING SOLUBILITY

• Nature of solute & Solvent
  • Solubility: refers to the maximum amount of solute expressed in grams that can be dissolved in 100g of water
  • Miscibility: ability of one substance to mix with another substance
• Temperature
  • ⬆ temperature: ⬆ solubility of a solid in a liquid
  • ⬆ temperature: ⬇ solubility of a gas in a liquid
  • Exothermic: ⬆ temperature: ⬆ solubility
  • Endothermic: ⬆ temperature: ⬇ solubility of a gas in a liquid
• Pressure
  • Henry’s law of gas solubility state that at constant temperature the solubility of a gas in a liquid is directly proportional to the pressure of the gas above the liquid
  • ⬆ pressure: ⬆ solubility of a gas
• Particle size
  • ⬆ Solubility: ⬇ Particle size
• Presence of salts
  • Salting-out: presence of salt decreases solubility.
  • Salting-in: presence of salt increases solubility.

TYPES OF SOLUTIONS

SOLUBILITY PRODUCT CONSTANT (KSP)

• Refers to the limit before saturation
  • Q < KSP: Unsaturated solution
  • Q = KSP: Saturated solution
  • Q > KSP: Supersaturated solution

IV. ACID-BASE EQUILIBRIA AND BUFFER

ACIDS

• Sour taste
• Litmus paper: Blue to red
• Strong Acids: HCBNIPS
  • HCl
  • HBr
  • HNO_3
  • HI
  • HClO_3
  • H2SO4

BASES

• Bitter taste
• Litmus paper: Red to blue
• Strong Bases:
  • OH of Group I and Group II Elements

SALTS

• Formed from neutralization reaction
• Types of salts:
  • Strong Acid + Strong Base = Neutral Salt
  • Strong Acid + Weak Base = Acidic Salt
  • Weak Acid + Strong Base = Basic Salt
  • Weak Acid + Weak Base = No Salt

THEORIES

IONIC EQUILIBRIA

• PH
  • Measured through Sorensen’s Scale
    • pH = 7: Neutral
    • pH>7: Basic
    • pH<7: Acidic
  • General Formula:
    • pH + pOH = 14
  • Dissociation constant:
    • Acid dissociation constant: High Ka = Acidic
    • Base dissociation constant: High Kb = Basic
• Ionic equilibria:
  • Strong Acid – Base: completely dissociate
  • Weak Acid – Base: does not completely dissociate
  • Non-electrolytes: does not dissociate

BUFFERS

• BUFFER
  • A combination of weak acid and its conjugate base or vice versa that resists changes in pH upon the addition of small quantities of acid or alkali.
• BUFFER CAPACITY
  • The magnitude of the resistance of a buffer to pH changes.

V. CHEMICAL REACTIONS AND ELECTROCHEMISTRY

TYPES OF CHEMICAL REACTIONS

• Composition/Direct Union:
  • A + B = AB
• Decomposition/Analysis:
  • AB = A + B
• Single Replacement
  • Depends on the activity series
    • Li – most reactive
    • Au- least reactive
• Double displacement
  • AB + CD = AD + BC

REDOX REACTION

• REDUCTION (VD GEROA)
  • Valence decreases and Oxidation State decreases
  • Gains electrons
  • Oxidizing agent
  • Addition of Hydrogen (Hydrogenation)
• OXIDATION (VI LEORA)
  • Valence increases and Oxidation State increases
  • Lose electrons
  • Reducing agent
  • Removal of Hydrogen (Dehydrogenation)
• ELECTROCHEMISTRY
  • RED CAT ELECT IN
    • Reduction occurs in the Cathode where Electrons Get In

VI. ATOMS AND SUBATOMIC PARTICLES

ATOM

• The basic unit of matter
• It is composed of the following
  • Nucleon:
    • Proton – positive charge
    • Neutron – neutral charge
  • Electrons:
    • 1,836x lighter than proton
    • Weight is negligible
    • First sub-atomic particle discovered

CONTRIBUTORS

NUCLIDE SYMBOL

• Mass number – Protons + Neutrons
• Atomic number = # of Protons
• #Protons = #Electrons in a neutral state

ISOTOPE, ISOBAR, ISOTONE

• ISOTOPE
  • Same element, same protons, different # of neutrons
  • Isotopes of Hydrogen
    • Protium – most abundant
    • Deuterium – Heavy hydrogen, used in NMR studies
    • Tritium – least abundant, only radioactive isotope of hydrogen
• ISOTONE
  • Different elements, same # of neutrons
• ISOBAR
  • Different elements, same mass #

VII. ELECTRONIC PRINCIPLES AND QUANTUM NUMBER

ELECTRONIC PRINCIPLE

• AUFBAU’S BUILDING UP PRINCIPLE
  • Electrons are first placed on the subshell with lower energy level.
• HUND’S RULE OF PAIRING
  • Electron orbitals are filled up singly.
    • Paramagnetism: attracted to a magnet
    • Diamagnetism: repelled by a magnet
• HEISENBERG’S UNCERTAINTY PRINCIPLE
  • It is impossible to determine simultaneously the momentum and position of an electron
• PAULI’S EXCLUSION PRINCIPLE
  • No 2 electrons can have the same exact set of quantum numbers

QUANTUM NUMBERS

• PRINCIPAL QUANTUM NUMBER (n)
  • Determines the main energy level/electron shell and size of the orbital
  • Values: Positive integers (1,2,3)
• AZIMUTHAL/ANGULAR QUANTUM NUMBER (l)
  • Determines the shape and subshell
  • Values: 0 to (n-1)

• MAGNETIC QUANTUM NUMBER (m_l)
  • Determines the special orientation
  • Values: -1 to +1
• SPIN QUANTUM NUMBER (m_s)
  • Describes the spin or rotation
  • Values:
    • +1/2 = clockwise
    • -1/2 = counterclockwise

VIII. PERIODIC TABLE OF ELEMENTS

PARTS OF THE PERIODIC TABLE

• There are 118 elements
• Periods (Horizontal)
  • 7 periods
  • Left to right rows
• Group/Family (Vertical)
  • Columns
  • Top to bottom
• Bridge element
  • Closely resembles the second member of an adjacent group to the right

GROUPS

• Group A
  • Also known as representative elements
  • S and P block
  • Known oxidation state
• Group B
  • Also known as transition elements
  • D and F block
  • Transition oxidation state

PERIODIC TRENDS

• PERIODIC TRENDS
  • Property that increases from right to left, top to bottom
    • Metallic property
    • Atomic Radius – ½ the distance between 2 nuclei
  • Property that increases from left to right; bottom to top
    • Ionization Energy – amount of energy required to remove an electron from a neutral atom to convert it to a positively charged ion
    • Non-metallic property
    • Electronegativity – tendency to attract electrons to itself
    • Electron Affinity – amount of energy released when a neutral atom accepts an electron to convert it to a negatively charged ion

METALS VS NON-METALS

METALLOIDS

• Intermediate between metals and non-metals
• Si Ge Po Sb ni Ar Te Bo
  • Si
  • Ge
  • Po
  • Sb
  • As
  • Te
  • B

IX. THERMODYNAMICS

LAWS OF THERMODYNAMICS

• ZEROTH LAW
  • states that if two bodies are each in thermal equilibrium with some third body, then they are also in equilibrium with each other
• FIRST LAW
  • Law of conservation of energy
    • Energy is neither created nor destroyed but is transformed
• SECOND LAW
  • Spontaneous Law
    • Degree of randomness/entropy
• THIRD LAW
  • The entropy of a pure crystalline solid at zero kelvin is zero

TYPES OF THERMODYNAMICS

GIBBS FREE ENERGY

• Measure of spontaneity
• General formula: G = \triangle H – T \triangle S
  • G = 0: Equilibrium
  • G > 0: Non-spontaneous/Endergonic
  • G < 0: Spontaneous/Exergonic

GIBBS FREE ENERGY SUMMARY TABLE

X. GAS LAWS FORMULA LIST

• Below is the list of formula for the different gas laws.

XI. NUCLEAR CHEMISTRY

NUCLEAR CHANGE

• Change in structure, properties, composition of the nucleus of an atom resulting in the transformation of the element into another element.
  • Nuclear Fission – splitting of heavy atom
  • Nuclear Fusion – union of 2 light atoms

NUCLEAR CHEMISTRY UNITS

• Units of radioactivity
  • Non-Si: Ci – 3.7 x 10^{10} decays/sec
  • SI: Bq – 1 decay/sec

TYPES OF RADIATION

XII. CHEMICAL KINETICS

CHEMICAL KINETICS

• Refers to the study of rate (or speed) of chemical reactions
• aA + bB ➡ cC + dD
• Chemical reaction – finished reaction
  • Shift to the right – forward reaction
  • Shift to the left – backward reaction

FACTORS AFFECTING RATE OF REACTION

• COLLISION THEORY
  • ⬆ rate = ⬆ collisions
  • Activation energy (E_a) – must be overcome for the action to proceed
• TRANSITION STATE THEORY
  • Reactants will undergo a transition state in order to produce products
• OTHER FACTORS
  • Particle size
  • Concentration
  • Nature of reactant
  • Temperature
  • Catalyst
    • Speeds up the reaction or lowers activation energy

CHEMICAL EQUILIBRIUM

• LE CHATELIER’S PRINCIPLE
  • If a stress is applied or placed on an equilibrium system, the system will respond by altering the equilibrium in such a way as to minimize stress.
• FACTORS AFFECTING EQUILIBRIUM
  • Concentration
    • Add A or B = forward reaction
    • Add C or D = backward reaction
  • Pressure
    • ⬆ Pressure = ⬇ Volume = less no. of moles
    • ⬇ Pressure = ⬆ Volume = more no. of moles
  • Temperature
    • Reactant – Endothermic
    • Product – Exothermic
• LAW OF MASS ACTION
  • States that the rate of the reaction Is proportional to the product of the concentrate of the reactants to the power of its coefficient in a balanced equation.