Unit 1: Sig Figs, Dimensional Analysis

• Sig Figs

  • If there is a decimal present, you start on the left side of the number with the first non-zero and count to the end. (Example: 1.250 has 4 sig figs, because 0 marks precision)

  • If a decimal isn’t present, start on the Atlantic side (right side) and count to the end. (Example: 12500, there are only 3 sig figs: 1, 2, and 5 because the two zeros’ do not mark precision)

  • When adding or subtracting decimals, the answer should be rounded to contain the least number of decimals based on the numbers added (Example: 123.25 has 2 decimals places + 46.0 only 1 decimals place + 86.257 has 3 decimals places, when added together, you get 255.607, but because 46.0 had one decimal place it should be rounded to 255.6)

  • When multiplying or dividing, reduce the number to the least amount of sig figs (Example: 23. 0 has 3 sig figs times 432 also has 3 sig figs x 19 has 2 sig figs, the answer is 188,784, but because 19 only has two sig figs the number should be rounded to 190,000 which has only 2 sig figs)

• Scientific Notation

  • Scientific Notation looks like N x 10^m

  • N is a number between 1 and 9.99 and cannot be any higher or lower

  • M is any exponent that is a whole number.

  • If there is a negative exponent, then you should move the decimal to the left (Example: 3.45 × 10^-3, becomes 0.00345)

  • If there is a positive exponent you move the decimal to the right (Example: 3.45 × 10³ is 3450

• Dimensional Anaylsis

  • Use Table C

  • To convert you first write your term to be converted,then multiply by the conversion factor, and solve. (Example: When converting 50 miles to kilometers you would write the term to be converted which is 50, then the conversion factor which is 1.60934km/1 miles because that is what you are converting, then multiply and you get that 50 miles is 80.467 kilometers)

Unit 2: Matter & Energy

• Matter

  • Anything that has mass and volume

  • Matter is made of 2 groups pure matter which are substances and impure matter which are mixtures

• Pure Matter

  • Pure substances are composed of atoms with the same atomic number (Example: O2 is a pure substance but H2O isn’t because H and O have different atomic numbers)

  • Any pure substance CAN NOT be broken down into simpler substances

  • Pure matter includes all elements on the periodic table

  • Almost all elements are monotomic exceot for some diatomic, Have No Fear of Ice Cold Beer (H2, N2, O2, I2, Cl2 and Br2)

  • Compounds that are only CHEMICALLY COMBINED, and have the same traits from sample to to sample (Example: Water is chemically combined and has the same traits in any cup)

  • Properties of compounds are different from their respective elements (Example: NaCl is table salt, but Na is a toxic metal and Cl is a toxic gas)

• Impure Matter

  • Mixtures

  • Mixtures are two or more substances that are PHYSICALLY COMBINED

  • Mixtures always retain the properties of their original components

  • They do not have the same compsition, they vary from sample to sample

  • There are two types of mixtures Homogenous and Heterogenous

  • Homogenous is evenly mixed are are made of water (Example: Salt Water)

  • Heterogenous is not evenly mized and no two samples will be the same (Example: Cereal)

• Physical & Chemical Properties of a Substance

  • Physical

  • There is no change in the composition

  • Differences in properties such as density, melting point, boiling point, atomic radius make it easy to seperate the substance

  • Physical changes allow use to change the substance without any chemical change (Example: Phase changes, dissolving it into individual ions and changing the size)

  • Words that are associated with physical change are color, texture, density, melting and boiling point, conductivity, luster and hardness

  • Chemical

  • Words that accosiate with chemical properties include burning, decomposing, spoils, reacts with, combines, rusts, reactivity

  • A change in the chemical compisition of a substance

  • It produces a new substance with different properties

• Intensive vs. Extensive Properties

  • Intensive properties do not depend on the amount of matter (Example: Temperature, Boiling Point, Concentration and luster)

  • Extensibe Properties depend on how much matter a sample contains (Example: weight, length, volume, and entropy)

• Seperation of Mixtures

  • Seperation with a magnet used to draw out a metal (Physical - Heterogenous)

  • Filtration, seperates by particle size (Physical - Heterogenous)

  • Seperation based on evaporation (Physical - Homogenous)

  • Distillation, used to seperate liquids or liquids + solids can be seperated by boiling point (Physical - Homogenous)

  • Chromotography, Mixture is dissolved into a solvent and compnents move through the phaze at different speeds (Physical - Homogenous)

  • Centrifuge, speration based on density of objects (Physical- Homogenous)

• Energy & Types of Energy

  • Energy is the ability to do work

  • Physical and chemical changes are changes in energy

  • There are two forms of energy: Mechanical and Non Mechanical

  • Mechanical energy has two components, potential energy, which is STORED energy that does not move and kinetic energy, that describes the MOTION of the molecules

  • Non mechanical energy is made up of five components, chemical energy, light energy, electrical energy, and atomic/nuclear energy

  • Thermal Energy is another important form of Non-Mechanical Energy

  • In heat energy heat flows from hot to cold until they are both the same temp, it is also the energy associated eith the random movment of atoms, molecules and the mass of matter, the unit is Joules (j)

  • Temperature is the measure of the average amount of molecules moving in a substance (Kinetic Energy), the unit is Kelvin or Celsius

  • The higher temperature the more kinetic energy and the lower temperature is the lower the kinetic energy

• Phase Changes

  • Solids have a definite shape, have a definite volume and their particles vibrate back and forth

  • Liquids take teh shape of the container, have a definite volume and the particles glide against each other

  • Gases have no definite shape or volume and the molecules are far apart, and move in straight lines.

  • From solid to liquid melting occurs, from liquid to gas evaporation occurs and from gas to solid deposition occurs

  • From gas to liquid condensation occurs, from liquid to solid freezing occurs and from solid ro gas sublimation occurs

• Types of Reactions for Phase Changes

  • Energy is either released or absorbed during a physical or chemical change

  • Exothermic and Endothermic

  • In Exothermic energy is released, this means chemical energy to heat energy and is used in deposition (g→ s) condensation (g→ l) and freezing (l → s)

  • In endothermic energy is absorbed, this means heat energy to chemical energy, this is used in sublimation (s → g) and (l → g)

• Diagrams

  • In a heating curve solid to liquid to gas is an endothermic process which means heat is absorbed

  • In a cooling curve from gas to liquid to solid is an exothermic process which means heat is released

  • In a heating curve kinetic energy increases and decreases in a cooling curve

  • Intermolecular forces between the molecules decreases in a heating curve is decreasing and increasing in a cooling curve)

  • Melting point is when the substance starts to melt

  • Freezing point when the force of attraction increases and the molecules move closer and liquid becomes a solid

  • Boiling point is when the force of attraction decreases and liquid turns into a gas

• Heat Equations

  • Heat of fusion is the amount of energy needed to convert from solid to liquid or vice versa

  • Heat of vaporization is the amount of energy needed to convert from liquid to gas or vice versa

  • The specific heat capacity of water is how much energy one gram of water needs to absorb to increase the temperature by one degree (4.18 j/g x c)

  • Table B

  • Heat of reaction equation is used IN a phase

  • Heat of fusion is used WHEN phase changing (s → l or l → s)

  • Heat of vaporization is used when phase changing (l→ g or g → l)

Unit 3: Atomic Theory

• Discoveries of the Atom

  • Dalton found that the basic unit of matter is an atom, they aren't divisible and all atoms in an element are identical, atoms of different elements are different.

  • Rutherford discovered the nucleus by using charged alpha particles and shooting them at gold foil, it deflected off something, showing something positively charged.

  • Thomson discovered the electrons using a cathode ray, he also discovered that an atom was a ball of positive charge that had negative electrons floating in it.

  • Gold Foil experiment found that the atom is mostly empty, and has a small dense positively charged nucleus in the center.

  • Bohr model was a model that had a small nucleus, had orbitals that contain electrons and each electrons has to have the right amount of energy in order to be placed in the orbital

  • Wave Mechanical Model found that electrons have energy and act as waves and particles, they’re found in clouds which have orbitals, where electrons are based on the energy they have.

• Parts of the Atom

  • Nucleus is the center of the atom it contains protons which have positive charge and neutrons which have no charge

  • Electrons orbit the proton in outer shells, they are negatively charged

  • Protons have a mass of 1u, charge of +1 and symbol of 1, 1 H

  • Neutrons have a mass of 1u, no charge, and a symbol of 1, 0 n

  • Electrons have a much smaller mass then protons and neutrons, have a charge of -1 and have a symbol of 0, -1e

  • On the Periodic table the number of protons is equal to the number of electrons

  • The atomic number is the number of protons in the nucleus

  • Mass is found by adding the number of protons by the number of neutrons, this is the number found on the periodic table rounded to the nearest whole number.

• Isotopes & Ions

  • Isotopes

  • Different atoms of the same element that have different mass numbers, they only have different amount of neutrons

  • Ions

  • These are charged atoms that either have positive or negative charges.

  • If they are positively charged they have more protons then electrons

  • If they are negatively charged they have more electrons then protons

  • The amount of protons STAY THE SAME all that changes is the amount of electrons, they either get less or get more.

  • Cations Vs. Anions

  • Cations are positively charged, they have fewer electrons then protons (Example: N+3, the +3 indicates that there are 3 fewer electrons then protons, so N+3 would have 7 protons and 4 electrons)

  • Anions are negatively charged and have more electrons then protons (Example: Cl-1, the -1 indicates that there is 1 more electron, so it would have 17 protons and 18 electrons)

• Atomic Mass & Electron Configuration

  • Atomic mass is the average mass of all naturally occurring isotopes of an element

  • Atomic Mass =(mass of isotope 1 x abundance of isotope 1) + (mass of isotope 2 x abundance of isotope 2) + …. / 100

  • Principal Energy Levels (PEL’s)

  • Tells you have far the electron is from the nucleus

  • PEL 1 is the lowest energy and can hold 2 electrons

  • PEL 2 can hold 8 electrons

  • PEL 3 can hold 18 electrons

  • PEL 4 has the highest energy (of these 4 only) and holds 32 electrons

  • As PEL’s get higher they gold more energy (Example: PEL 5 has more than PEL 4)

  • If a PEL is occupied that means that it has electrons in it but it does not have the maximum amount of electrons (Example: Fluorine has a configuration of 2-7 and the first PEL is filled but the second is occupied cause it can hold one more electron)

  • If the PEL is filled, it has the maximum amount of electrons in it.

  • Electron Configuration

  • Shows how many electrons are in each PEL

  • Valence electrons are the electrons found in the outer shell, this is also the last number in electron configuration

  • Ground State Vs. Excited State

  • Electrons located in orbitals have the energy of that orbital, but they can move to other orbitals by absorbing energy and jumping around

  • The energy absorbed is absorbed in the form of heat or light

  • The ground state is the most stable, everything on the periodic table is in the ground state

  • Excited state is unstable, electrons jump up and leave a PEL partly empty

  • Atomic Emission Spectra

  • Electrons do not stay in the excited state for a long time, that would be very unstable

  • When electrons jump back down they give off energy in the form of color

  • Every element has a different emission spectra

Unit 4: The Periodic Table

• Arrangement

  • The table is organized in order of atomic number (Number of protons)

  • Periods are the rows in the periodic table

  • Period number is the number of the PEL shells and the energy level where valence electrons are located

  • When moving across the period add one electron to the VALENCE SHELL and one proton each time

  • Groups are the columns in the periodic table

  • Atoms in each group have the same amount of valence electrons

  • If the atoms are in the same group they have the same amount of valence electrons which also means they have similar physical and chemical properties

• Organization

  • Metals

  • Metals are on the left of the staircase

  • Metals are usually solid except for mercury which is a liquid

  • Metals are conductors of heat and electricity

  • They are shiny

  • They are bendable and ductile

  • They have a low ionization energy (energy required to be an ion) and low electronegativity (ability to accept electrons)

  • Metals become cations with small ionic radius’

  • Francium is the most reactive metal

  • Metals have fewer electrons in the VALENCE shell

  • Nonmetals

  • They are on the right of the staircase

  • Solids are iodine

  • Liquids are bromine

  • Gasses are all the diatomic except for Ice (Iodine)

  • They are poor conductors of heat and electricity, aren’t shiny and are brittle

  • They have high ionization and electronegativity

  • They become anions with large ionic radius’

  • Fluorine is the most reactive nonmetal

  • They have more electrons in the VALENCE shell

  • Metalloids

  • ON THE STAIRCASE

  • Have characteristics of metals and nonmetals

  • 6 elements: Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb) and Tellurium (Te)

  • BE, like SILICON, and GERMAN ARSENIC weapons and don’t let old ANTIMONY TELL-UR-MUM

• Transition Metals (3 - 12)

  • Multiple oxidation states

  • Form ions with color

  • They are the least reactive groups of metals

  • Most likely to be found combined

• Group 1 (Alkali Metals)

  • They have one valence electron

  • Very rarely found

  • They are soft and very reactive

  • They are the entire first column except for hydrogen

• Group 2 (Alkaline Earth Metals)

  • Two valence electrons

  • very rarely found

  • All elements in the second column

• Group 17 (Halogens)

  • The most reactive non metals

  • Only group to contain all 3 states of matter

  • They are all diatomics

• Group 18 (Noble Gases)

  • Unreactive

  • Have oxidation numbers of 0

  • Krypton and Xenon have oxidation states but rarely react

  • All are octets and have 8 electrons

• Allotropes

  • Elements in the same phase and can exist in two or more structures

  • Oxygen → Oxygen (O2) and Ozone (O3)

  • Carbon → Diamond, Graphite and Coal

• Trends

  • Ionization decreases down a group and increases across a period

  • Electronegativity decreases down the group and increases across a period

  • Atomic radius increases down the group and decreases across the period

  • Ionic Radius of metals are smaller than atomic radius and ionic Radius in nonmetals are greater than the atomic radius

Unit 5: Bonding

• Bonds

  • Bonds are forces of attraction between protons of one atom and another atwom

  • Only valence electrons

  • Result of a chemical reaction

  • Bonds are spontaneous and naturally occur

  • Energy is released in bond forming

  • Energy goes from high to low

  • Bond breaking is not spontaneous

  • Energy is absorbed in bond breaking

  • Goal is to have a complete outer shell

• Criss Cross Rule + Reverse Criss Cross Rule

  • Take the oxidation number of each element and criss Cross them

  • Do not bring down the charge, just the number (Example: Aluminum Oxide Al 3+ and O-2 becomes Al2O3

  • For the reverse Criss cross rule just switch the subscripts back to the oxidation states of each element

  • Metals form cations which means their oxidation states will be positive

  • Nonmetals form anions which means their oxidation states will be negative

• Ionic Bonds

  • Formed with a metal and nonmetal

  • Metal atoms transfer electrons to a nonmetal atom

  • They have a large electronegativity difference of 1.7 or higher

  • They are hard substances

  • They have high melting and boiling points

  • They can dissolve in water to form solutions and are good conductors of electricity in solutions

  • Dissolve in polar substances ONLY

• Covalent Bonds

  • A bond formed between two nonmetals

  • They SHARE electrons

  • The electronegativity difference is 1.6 or less

  • There are two types of bonds Polar and Nonpolar Covalent Bonds

• Polar Covalent Bonds

  • Have an electronegativity difference of 0.4 - 1.6

  • They form DIPOLES, which is one positive end and one negative end

  • They have UNEQUAL sharing of electrons

  • Have Asymmetrical shape

• Nonpolar Covalent

  • Have and electronegativity difference of o - 0.3

  • They have NO dipoles

  • They have EQUAL sharing

  • They are symmetrical

• Metallic Bonds

  • Found in metals

  • Have a sea of mobile electrons

  • Hard substances with HIGH melting and boiling points

• Criss Cross Rule With Polyatomics

  • Polyatomic ions are two or three elements combined as one ion

  • They act as one unit and DO NOT BREAK, when Criss crossing subscripts are written outside of the paranthesis

• Naming Binary Compounds

  • Metal + Nonmetal is written as the metal first and the nonmetal second, but the ending is changed to ide

  • Metal + Polyatomic ion is written as metal first, polyatomic second, and BOTH keep their names

  • Polyatomic + Nometal is written as the polyatomic first, nonmetal second and the ending is changed to ide

  • Polyatomic Ion (Cation) + (Anion) both keep their names

• Naming Covalent Compounds

  • First nonmetal (less electronegative) keeps its name

  • Second Nonmetal ( more electronegative) changes to IDE

  • Mono is not needed for the first nonmetal if there is only one but if there are two it needs a prefix.

  • Mono - 1, Di - 2, Tri -3, Tetra -4, Penta - 5, Hexa - 6, Hepta - 7, Octa - 8, Nona - 9, Deca - 10

• Transition Metal Naming

  • Uses Roman numerals to identify oxidation states

  • Metal keeps the name, nonmetal changes to ide and polyatomics keep the name

• Molecular Polarity

  • Asymmetrical Shape: Different charge distribution on all sides.

  • Polar Bonds: Electronegativity difference from 0.4 to 1.6; unequal sharing of electrons, leading to unequal charge distribution and formation of dipoles (e.g., Hydrogen Chloride, Ammonia, Water).

  • Nonpolar Molecules: Symmetrical shape; dipoles pull equally on all sides. Can have polar bonds but overall nonpolar due to charge distribution (e.g., Carbon Dioxide, Carbon Tetrachloride, Methane).

• Intermolecular Forces

  • Hydrogen Bonds, strong intermolecular forces responsible for the high boiling point of water formed when Hydrogen bonds with Fluorine, oxygen or nitrogen

  • Dipole, Diople, attraction between the positive and negative end of dipole

  • London Dispersion Forces, interactions between electron clouds, and weak forces between nonpolar molecules. The strength of the bond increases with size

    Physical Properties Dependent on Intermolecular Force Strength

    • Melting and boiling points.

    • Surface tension.

    • Viscosity.

    • Solubility.

    Gas Laws

    • Review of Matter: Standard Temperature and Pressure (STP) – 1 atm or 101.3 kPa; 273 K.

    • States of Matter: Solids (strong intermolecular forces), Liquids (strong intermolecular forces), Gases (weak intermolecular forces).

    Vapor Pressure & Table H

    • Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid. Increases with temperature.

    Combined Gas Law

    • Pressure and volume are inversely proportional (Boyle’s Law).

    • Pressure and temperature are directly proportional (Gay-Lussac’s Law).

    • Temperature and volume are directly proportional (Charles’s Law).

    Ideal Gas vs. Real Gas

    • Ideal gases follow gas laws perfectly, while real gases deviate at high pressures and low temperatures due to intermolecular forces.

    Unit 6: Chemical Reactions and Stoichiometry

    Types of Chemical Reactions

    1. Synthesis: Two or more simple substances combine to form a more complex substance (A+B→ABA + B \rightarrow ABA+B→AB).

    2. Decomposition: A complex substance breaks down into simpler substances (e.g., AB→A+BAB \rightarrow A + BAB→A+B).

    3. Single Replacement: One element replaces another in a compound (e.g., A+BC→AC+BA + BC \rightarrow AC + BA+BC→AC+B).

    4. Double Replacement: Exchange of ions between two compounds (e.g., AB+CD→AD+CBAB + CD \rightarrow AD + CBAB+CD→AD+CB).

    5. Combustion: A substance combines with oxygen, releasing energy (e.g., CxHy+O2→CO2+H2OC_xH_y + O_2 \rightarrow CO_2 + H_2OCx​Hy​+O2​→CO2​+H2​O).

    Balancing Chemical Equations

    • Ensure the number of atoms of each element is the same on both sides of the equation.

    • Follow the law of conservation of mass.

    Stoichiometry

    • Mole Ratios: Use coefficients from balanced equations to determine ratios.

    • Calculations:

      • Mass-to-Mass: Convert grams of a reactant to moles, use mole ratio to find moles of product, then convert to grams.

      • Volume of Gases: Use ideal gas law PV=nRTPV = nRTPV=nRT to relate volume, pressure, and temperature.

    Limiting Reactant

    • Determines the amount of product formed in a reaction; the reactant that runs out first.

    • Steps: Calculate moles of each reactant, use stoichiometry to determine which produces the least product.

    Percent Yield

    • Actual Yield: Measured amount obtained from a reaction.

    • Theoretical Yield: Maximum amount predicted by stoichiometry.

    • Percent Yield: (Actual YieldTheoretical Yield)×100\left( \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \right) \times 100(Theoretical YieldActual Yield​)×100.

    Unit 7: Solutions

    Nature of Solutions

    • Solvent: Substance that dissolves the solute.

    • Solute: Substance being dissolved.

    • Types of Solutions: Gaseous, liquid, solid.

    Concentration of Solutions

    • Molarity (M): Moles of solute per liter of solution.

    • Dilution: M1V1=M2V2M_1V_1 = M_2V_2M1​V1​=M2​V2​, where MMM is molarity and VVV is volume.

    Solubility

    • Factors Affecting Solubility:

      • Temperature: Generally increases solubility of solids, decreases for gases.

      • Pressure: Affects gas solubility (Henry’s Law).

    Colligative Properties

    • Depend on number of solute particles:

      • Boiling Point Elevation: Solution boils at a higher temperature than pure solvent.

      • Freezing Point Depression: Solution freezes at a lower temperature than pure solvent.

      • Vapor Pressure Lowering: Presence of solute lowers vapor pressure of solvent.

      • Osmotic Pressure: Pressure required to stop osmosis.

    Electrolytes and Nonelectrolytes

    • Electrolytes: Substances that conduct electricity when dissolved in water (e.g., salts).

    • Nonelectrolytes: Do not conduct electricity (e.g., sugar).

    Acids and Bases

    • Properties:

      • Acids: Sour taste, conduct electricity, react with metals.

      • Bases: Bitter taste, slippery feel, conduct electricity.

    • pH Scale: Measures acidity/basicity; pH<7\text{pH} < 7pH<7 is acidic, pH>7\text{pH} > 7pH>7 is basic, pH=7\text{pH} = 7pH=7 is neutral.

    • Neutralization Reaction: Acid and base react to form water and salt (e.g., HCl+NaOH→NaCl+H2O\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}HCl+NaOH→NaCl+H2​O).

Unit 8: Solutions

Solutions and Solubility

• Solutions: Homogeneous mixtures of two or more substances. The substance in the greatest amount is the solvent, and the substance in lesser amounts is the solute.

• Types of Solutions: Solid, liquid, and gas solutions depending on the states of the solute and solvent.

• Solubility: The maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature. Factors affecting solubility include temperature, pressure, and the nature of the solute and solvent.

• Saturation:

  • Unsaturated Solution: Contains less solute than the maximum amount that can be dissolved.

  • Saturated Solution: Contains the maximum amount of solute that can dissolve.

  • Supersaturated Solution: Contains more solute than can theoretically dissolve at a given temperature; achieved by cooling a saturated solution.

Concentration

• Concentration: The amount of solute dissolved in a given quantity of solvent. Common units include molarity (M), molality (m), and percent composition.

• Molarity (M): Moles of solute per liter of solution.

• Molality (m): Moles of solute per kilogram of solvent.

• Percent Composition: Mass of solute per mass of solution, multiplied by 100%.

Factors Affecting Solubility

• Temperature:

  • Solubility of solids typically increases with temperature.

  • Solubility of gases decreases with increasing temperature.

• Pressure:

  • Affects the solubility of gases significantly (Henry's Law).

  • Increased pressure increases gas solubility in liquids.

• Nature of Solute and Solvent: "Like dissolves like" principle; polar solvents dissolve polar solutes, and nonpolar solvents dissolve nonpolar solutes.

Colligative Properties

• Properties that depend on the number of solute particles in a solution, not the type of particle. Includes boiling point elevation, freezing point depression, vapor pressure lowering, and osmotic pressure.

• Boiling Point Elevation: Addition of solute raises the boiling point of the solvent.

• Freezing Point Depression: Addition of solute lowers the freezing point of the solvent.

• Vapor Pressure Lowering: Addition of solute lowers the vapor pressure of the solvent.

• Osmotic Pressure: Pressure required to stop the flow of solvent into the solution through a semipermeable membrane.

Unit 9: Kinetics and Equilibrium

Reaction Rates

• Reaction Rate: The change in concentration of reactants or products per unit time.

• Factors Affecting Reaction Rates:

  • Concentration: Higher concentration of reactants generally increases the reaction rate.

  • Temperature: Higher temperature increases reaction rate by providing more energy to reactant molecules.

  • Surface Area: Greater surface area of a solid reactant increases reaction rate.

  • Catalysts: Substances that increase reaction rate without being consumed by lowering the activation energy.

  • Nature of Reactants: Different substances react at different rates based on their chemical nature.

Chemical Equilibrium

• Dynamic Equilibrium: When the rate of the forward reaction equals the rate of the reverse reaction in a closed system.

• Equilibrium Constant (K): The ratio of product concentrations to reactant concentrations at equilibrium, with each concentration raised to the power of its coefficient in the balanced equation.

• Le Chatelier's Principle: If a system at equilibrium is disturbed, it will shift to counteract the disturbance and restore a new equilibrium. Factors include changes in concentration, temperature, and pressure.

Equilibrium Calculations

• ICE Tables: Used to calculate changes in concentration and equilibrium concentrations (Initial, Change, Equilibrium).

• Reaction Quotient (Q): Calculated like the equilibrium constant but with initial concentrations; used to predict the direction the reaction will shift to reach equilibrium.

Unit 10: Acids and Bases

Acid-Base Theories

• Arrhenius Definition:

  • Acid: Produces H⁺ ions in solution.

  • Base: Produces OH⁻ ions in solution.

• Bronsted-Lowry Definition:

  • Acid: Proton (H⁺) donor.

  • Base: Proton (H⁺) acceptor.

• Lewis Definition:

  • Acid: Electron pair acceptor.

  • Base: Electron pair donor.

Properties and Strengths

• Properties of Acids: Sour taste, turn blue litmus red, react with metals to produce H₂ gas, conduct electricity in solution.

• Properties of Bases: Bitter taste, slippery feel, turn red litmus blue, conduct electricity in solution.

• Strong vs. Weak Acids/Bases: Strong acids/bases dissociate completely in water, while weak acids/bases dissociate partially.

pH and pOH

• pH: Measure of hydrogen ion concentration; pH = -log[H⁺]. Scale ranges from 0 (very acidic) to 14 (very basic), with 7 being neutral.

• pOH: Measure of hydroxide ion concentration; pOH = -log[OH⁻]. pH + pOH = 14.

• Indicators: Substances that change color at specific pH values, used to determine the pH of a solution.

Titration

• Titration: Technique used to determine the concentration of an unknown solution by reacting it with a solution of known concentration.

• Equivalence Point: Point in titration where the amount of acid equals the amount of base.

• Endpoint: Point in titration where the indicator changes color, ideally close to the equivalence point.

Unit 11: Oxidation-Reduction

Oxidation States

• Oxidation States: Indicates the number of electrons lost or gained by an atom in a compound. Rules include:

  • Free elements have an oxidation state of zero.

  • For ions, the oxidation state is equal to the charge.

  • Specific elements have consistent oxidation states (e.g., Group 1 metals are +1, Group 2 metals are +2, Fluorine is always -1).

Redox Reactions

• Oxidation: Loss of electrons; increase in oxidation state.

• Reduction: Gain of electrons; decrease in oxidation state.

• Oxidizing Agent: Substance that gets reduced and causes oxidation.

• Reducing Agent: Substance that gets oxidized and causes reduction.

Balancing Redox Equations

• Half-Reaction Method: Separates the oxidation and reduction processes, balances each half-reaction for mass and charge, and then combines them to form the balanced overall reaction.

• Steps:

  1. Write the oxidation and reduction half-reactions.

  2. Balance atoms other than H and O.

  3. Balance O atoms by adding H₂O.

  4. Balance H atoms by adding H⁺ (in acidic solution) or OH⁻ (in basic solution).

  5. Balance charge by adding electrons.

  6. Multiply half-reactions by appropriate coefficients to equalize the number of electrons transferred.

  7. Add the half-reactions and simplify.

Unit 12: Organic Chemistry

Overview

Unit 12 focuses on the study of carbon and carbon-based compounds. Carbon's unique ability to form four covalent bonds allows for a vast array of structures, including chains, branches, rings, and networks, making organic chemistry a diverse and essential field of study.

Characteristics of Organic Compounds

1. Non-Polar Nature:

   • Organic compounds are generally non-polar, meaning they are soluble in non-polar substances.

2. Non-Electrolytes:

   • These compounds typically do not conduct electric currents.

3. Low Melting and Boiling Points:

   • Organic compounds usually have low melting and boiling points.

4. Reaction Rates:

   • Organic reactions tend to be slow and require high activation energies.

5. Solubility:

   • Organic compounds are generally insoluble in water.

6. Boiling and Melting Points:

   • As the number of carbon atoms in a compound increases, so do the boiling and melting points due to increased intermolecular forces.

Structural Formulas and Isomers

• Structural Formula:

  • Shows the bonding patterns and shapes of molecules.

• Isomers:

  • Compounds with the same molecular formula but different structural formulas, leading to different physical and chemical properties.

  • Example: Acetone (C3H6O) and Propanal (C3H6O) are isomers with distinct properties.

Saturated vs. Unsaturated Compounds

• Saturated Compounds:

  • Contain only single bonds between carbon atoms.

• Unsaturated Compounds:

  • Contain at least one double or triple bond between carbon atoms.

Homologous Series (Table Q)

• Provides formulas for three sets of hydrocarbons (organic compounds containing only carbon and hydrogen):

  1. Alkanes: Saturated hydrocarbons with single bonds.

  2. Alkenes: Unsaturated hydrocarbons with at least one double bond.

  3. Alkynes: Unsaturated hydrocarbons with at least one triple bond.

Bonding Patterns

• Carbon: 4 valence electrons, forms 4 bonds.

• Oxygen: 2 bonds.

• Nitrogen: 3 bonds.

• Hydrogen: 1 bond.

• Halogens (Group 17): 1 bond.

Nomenclature and Prefixes

• Prefixes indicate the number of carbon atoms in the organic compound and are used for naming branches of carbon chains.

Functional Groups and Their Reactions

• Functional Groups:

  • Specific groups of atoms within molecules that determine the characteristic chemical reactions of those molecules.

• Common Functional Groups:

  1. Hydroxyl Group (-OH): Found in alcohols.

  2. Carbonyl Group (C=O): Found in aldehydes and ketones.

  3. Carboxyl Group (-COOH): Found in carboxylic acids.

  4. Amino Group (-NH2): Found in amines and amino acids.

  5. Halides (F, Cl, Br, I): Found in halogenated compounds.

  6. Ethers (R-O-R'): Compounds with an oxygen atom connected to two alkyl or aryl groups.

  7. Esters (R-COO-R'): Derived from carboxylic acids and alcohols.

• Reactions Involving Functional Groups:

  1. Substitution Reactions: Atoms in a molecule are replaced by other atoms or groups.

  2. Addition Reactions: Atoms or groups are added to unsaturated molecules (alkenes and alkynes).

  3. Elimination Reactions: Atoms or groups are removed from a molecule, often resulting in the formation of a double or triple bond.

  4. Condensation Reactions: Two molecules combine with the loss of a small molecule, often water.

  5. Oxidation-Reduction Reactions: Involves the transfer of electrons between molecules, affecting the oxidation states of atoms.

Organic Synthesis and Reactions

• Polymerization:

  • The process of reacting monomer molecules together to form polymer chains or three-dimensional networks.

  • Addition Polymerization: Monomers add to each other without the loss of any atoms.

  • Condensation Polymerization: Monomers join with the loss of a small molecule such as water.

• Combustion:

  • Organic compounds react with oxygen to produce carbon dioxide, water, and energy.

  • Complete Combustion: Produces CO2 and H2O.

  • Incomplete Combustion: Produces CO, C (soot), and H2O.

Environmental Impact and Uses

• Uses of Organic Compounds:

  • Pharmaceuticals, plastics, fuels, dyes, and many other materials.

• Environmental Impact:

  • Organic compounds can have significant environmental impacts, including pollution and contribution to climate change through the release of greenhouse gases.

This detailed summary provides an in-depth look at Unit 12, covering all essential topics related to organic chemistry, including characteristics, structures, functional groups, reactions, and their environmental implications.