Written Report (Inorganic Chem ).pdf

INORGANIC CHEMISTRY Chemical Change of Matter - occurs when a substance undergoes a chemical reaction and forms Matter - anything that takes up space and can be one or more new substances with diffe...

INORGANIC CHEMISTRY Chemical Change of Matter - occurs when a substance undergoes a chemical reaction and forms Matter - anything that takes up space and can be one or more new substances with different chemical weighed. In other words, matter has volume and properties. mass. THE ATOMIC THEORY States of Matter A scientific theory that explains the nature of matter by stating that all matter is composed of discrete Solid - particles are tightly packed with minimal units called atoms. movement, resulting in low kinetic energy. Solids have a definite shape, mass, and volume, and do not The development of atomic theory has evolved over conform to the shape of their container. time as scientific understanding has progressed. Liquid - particles are more loosely packed than in solids, allowing them to flow around each other. This Atomic Models gives liquids an indefinite shape, enabling them to conform to the shape of their container. Solid Sphere (John Dalton, 1803) Atoms are invisible and atoms of Gasses - particles have a lot of space between them elements are identical. Compounds and high kinetic energy. Gasses have no definite are combinations of atoms. shape or volume and will spread out indefinitely if unconfined. Plum Pudding / Thomson Model (J.J. Thomson, Plasma - is likely the most common state of matter 1904) in the universe, as seen in stars like the sun. Plasma consists of highly charged particles with extremely Electrons are scattered throughout a high kinetic energy. spherical cloud of positively charged particles. Bose-Einstein Condensate - occurs at extremely low temperatures, near absolute zero. In this state, a group of atoms is cooled to such low temperatures Nuclear / Rutherford Model (Ernest Rutherford, that they occupy the same quantum state, behaving 1911) as a single quantum entity with unique properties. Atom is mainly empty space with a positively charged center. Electrons Classifications of Matter revolve a predictable path. Planetary / Bohr’s Model (Niels Bohr, 1913) Negatively charged electrons revolve around the positively charged nucleus at a fixed orbit. Properties of Matter Quantum (Erwin Schrodinger, 1926) Electrons are found in clouds of probability called orbitals. Exact location is impossible to determine. Still widely accepted as the most accurate model of the atom. Atomic Structure Physical Change of Matter - is a change in the form or physical properties of a substance, without a change in its chemical composition. Subatomi Discovered Electric Charge Mass c Particle  Electron J.J. −1.602 × 10−19 9.11 × 10−31 Henry Moseley → Determined that atomic number, (e-) Thomson not atomic weight, is the correct basis for ordering Proton Ernest 1.673 1.602 × 10−19 elements in the Periodic Table. (p+) Rutherford × 10−21 Neutron James 1.675 (no) Chadwick 0 × 10−21 Johann Döbereiner → Identified the Law of Triads for elements with similar properties, where the atomic weight of the middle elements of the triad was Atomic Symbol roughly the average of the other two elements. John Newlands → Proposed the Law of Octaves, observing that every eighth element had similar properties when elements were arranged by atomic weight. Lothar Meyer → Independently developed a periodic table similar to Mendeleev's, recognizing the periodic trend between atomic volume and atomic weight. Robert Boyle → Defined elements as pure Number of protons = Atomic number substances that cannot be broken down into simpler Number of electrons = Number of protons substances, laying the groundwork for modern Number of neutrons = Mass number - Atomic chemistry. number Antoine Lavoisier → Known as the "Father of Atomic Properties Modern Chemistry", he developed a modern system Metallic Property - ability of naming chemical substances, identified and of an atom to donate an named oxygen and hydrogen, and established the electron Law of Conservation of Mass. Atomic Size - average PARTS OF THE PERIODIC TABLE distance between the nucleus and the valence Groups → Vertical columns on the Periodic Table; electron elements in the same group share similar valence electrons and chemical characteristics. Reactivity – tendency of an atom to react Periods → Horizontal rows on the Periodic Table; elements in the same period have the same number Ionization Energy – of electron shells. energy required to remove an electron from an atom Valence Electrons - Determines how a particular atom reacts (Same number of valence electrons = Electron Affinity – change in energy when a Reacts in the same similar manner) gaseous atom/ion gains an electron Main Group Elements → Elements in Groups 1A to Electronegativity – ability of an atom to attract or 8A. Also known as representative elements. gain electrons Transition Elements → Elements in Groups 3 to THE PERIODIC TABLE OF ELEMENTS 12. Characterized by their ability to form multiple oxidation states and have colored compounds. SIGNIFICANT CONTRIBUTIONS Classifications of Elements: Dmitri Mendeleev → Started the development of the periodic table, arranging elements by increasing 1. Metals atomic mass which led to the modern Periodic Table Usually shiny, very dense, high structure. melting points. Can be stretched into thin wires Antoine Bequerel → First discovered radioactivity. (ductile), hammered into thin sheets Pierre and Marie Curie → Discovered the elements (malleable), and are excellent radium and polonium conductors of heat and electricity.  In chemical reactions, they readily Actinides → Radioactive elements, many of which lose electrons to form positive ions. are synthetic. Uranium and thorium are notable for All are solid at room temperature their use in nuclear energy. except for mercury. Significant Elements: 2. Nonmetals Brittle, dull, low melting points. Hydrogen → A nonmetal, diatomic gas (H₂) Poor conductors of heat and that can act like an alkali metal by losing one electricity. electron or like a halogen by gaining one In chemical reactions, they tend to electron. gain electrons, forming negative ions. Oxygen → A highly reactive non-metal, essential for combustion and respiration. 3. Metals (or Semimetals) Forms oxides with most elements and is the Exhibit both properties of metals and most abundant element in the Earth's crust. nonmetals Carbon → Can form stable bonds with itself Semiconductors: Can conduct and other elements, creating a vast array of electricity better than insulators but organic compounds. The backbone of not as well as conductors organic chemistry and life. Elemental Composition of the Earth and the Human Body: CHARACTERISTICS OF ELEMENTS Alkali Metals (Group 1) → Highly reactive metals due to their larger atomic radii and low ionization energies. Soft, silvery, and can be easily cut. Alkaline Earth Metals (Group 2) → Reactive metals, but less so than alkali metals. They commonly form compounds like oxides and hydroxides. Transition Metals (Groups 3-12) → Characterized by their ability to form complex ions and are effective catalysts. Halogens (Group 7A/17) → Highly reactive nonmetals, known for forming salts when reacting PERIODIC TRENDS IN ATOMIC PROPERTIES with metals. Unique group because it includes all states of matter at room temperature and notable for Atomic Radius → Increases as you move down, containing 4 out of seven diatomic elements (F2, Cl2, Decreases as you move across (left to right). Br2, I2) Atomic size increases down a group due to added Noble Gases (Group 8A/18) → Inert, colorless, and electron shells, expanding the electron cloud. odorless gasses with a complete valence electron Across a period, it decreases as the increased shell, making them extremely stable and mostly non- number of protons in the nucleus exerts a stronger reactive. They are used in lighting and as inert pull on the electrons, reducing the atomic radius. atmospheres in chemical reactions. Lanthanides → Rare earth metals, known for their magnetic and phosphorescent properties. These elements are used in electronics, lasers, and strong magnets.  free element and reacts explosively with many substances. Electronegativity → Increases as you move up and TYPES AND NAMING OF COMPOUNDS across (left to right) Nomenclature Prefixes Fluorine remains the most electronegative element. Noble gasses, however, have an electronegativity Number Prefix Number Prefix rating of 0 due to their inherent stability, preventing them from forming bonds with other atoms. 1 mono- 6 hexa- 2 di- 7 hepta- 3 tri- 8 octa- 4 tetra- 9 nona- 5 penta- 10 deca- A. Ionic Compounds → Formed by the transfer of Ionization Energy → Increases as you move up electrons between metals and nonmetals. and across (left to right) Steps for Naming Ionic Compounds (Monatomic Depends on how tightly electrons are held by the Ions): nucleus; electrons closer to the nucleus are more tightly bound and require more energy to be Example: MgBr2 removed. 1. Name of Metal Ion: The ion’s name (Mg2+) Electron Affinity → Increases as you move up and is given in the periodic table as magnesium. across (left to right) 2. Name the non-metal ion by ending the element name with the suffix “ide”: The Electron affinity generally becomes more negative nonmetal ion is Br (Bromine). By changing across a period as atoms more readily accept the name to end with the suffix “ide”, it gives electrons to complete their valence shell, and “bromide”. becomes less negative down a group due to reduced 3. Write the name of the compound: nuclear attraction on added electrons. Magnesium bromide Metallic Character → Increases as you move Steps for Naming Ionic Compounds (Transition down, Decreases as you move across (left to right) Metal + Nonmetal): Refers to how easily an element can lose an Example: FeCl3 electron. Cesium is the most reactive metal, reacting explosively with water and igniting spontaneously in 1. Identify the Transition Metal and air. Francium, located below cesium in the alkali Nonmetal: Determine which transition metal metal group, is so rare that its properties are largely (Fe) and nonmetal (Cl) are present in the unobserved. compound. 2. Determine the Charges of the Ions: Nonmetallic Character → Increases as you move Transition metals can form cations with up and across (left to right) different charges. Find the charge of the transition metal ion (3+) from the formula of Refers to the tendency of an element to gain the compound and the charge of the electrons in chemical reactions. Fluorine is the most nonmetal ion (1-). reactive nonmetal as it is not found in nature as a  3. Write the Metal Name: Start with the name Named by combining the names of the cation and of the transition metal which is “Iron”. the anion (e.g., KNO₃ is potassium nitrate). 4. Specify the Metal Ion Charge: Indicate the charge of the transition metal ion using a G. Organic Compounds → Compounds primarily Roman numeral in parentheses immediately made of carbon and hydrogen atoms, often after the metal name. In this case, it will give containing oxygen, nitrogen, and other elements. “Iron (III)” Named based on the number of carbon atoms and 5. Name the Nonmetal: Follow the metal name the types of bonds (e.g., CH₄ is methane, C₂H₄ is with the name of the nonmetal which is ethene). “Chloride” 6. Write the name of the compound: Iron (III) QUANTUM CHEMISTRY chloride Quantum Chemistry → a branch of chemistry that B. Covalent (Molecular) Compounds → Formed studies and explains the behavior of subatomic by the sharing of electrons between nonmetals. particles like electrons and light. Steps for Naming Covalent Compounds: IV.I ELECTRONIC STRUCTURE OF THE ATOM Example: SiCl4 Electromagnetic Radiation → has oscillating electric and magnetic fields in planes perpendicular 1. Name the first element: The first element is to each other in the direction of propagation. Si thus, silicon. 2. Name the second element by ending the element name with the suffix “ide”: The second element is Cl thus it gives chloride. 3. Add prefixes to the atom names to indicate the number of each atom in the compound: Silicon is one atom (no prefix) while chlorine has four atoms which gives “tetrachloride” 4. Write the name of the compound: Silicon tetrachloride C. Polyatomic Ions → Ions made up of multiple Characterization of Electromagnetic Waves atoms bonded covalently but carrying an overall · Wavelength (𝛌) → distance between charge. Common polyatomic ions include sulfate two consecutive peaks or troughs; in (SO₄²⁻) and ammonium (NH₄⁺). Compounds meters (m) containing these ions are named by combining the · Amplitude → distance from origin to names of the cation and polyatomic ion (e.g., crest Na₂SO₄ is sodium sulfate). · Frequency (f) → number of waves/cycles that pass a point per unit D. Acids → Compounds that release H⁺ ions in time; in per seconds (s-1) or Hertz (Hz) solution. · Speed (v) → v = 𝛌f; shows the inverse relationship of frequency and wavelength a. Binary acids (containing hydrogen and one 𝑚 → speed of light =3 × 10−8 𝑠 other element) are named with the "hydro-" 𝑚 → speed of sound =343 prefix and "-ic" suffix (e.g., HCl is 𝑠 hydrochloric acid). b. Oxyacids (containing hydrogen, oxygen, and Max Planck → postulated that light and other another element) are named based on the electromagnetic waves can be quantized and were polyatomic ion; "-ate" becomes "-ic acid" and emitted in terms of discrete packets of energy called "-ite" becomes "-ous acid" (e.g., H₂SO₄ is Quanta ℎ𝑣 sulfuric acid). →𝐸 = ℎ𝑓 = 𝜆 ; Planck’s Constant (h) = or CONST 06 (for Canon F-789SGA) E. Bases → Compounds that release OH⁻ ions in solution. Named by stating the metal cation followed Rydberg’s Equation → use to quantify the by "hydroxide" (e.g., NaOH is sodium hydroxide). wavelength when the atom transitions from the excited phase to ground phase after exposure to F. Salts → Ionic compounds formed from the heat/flame. neutralization reaction between an acid and a base.  1 1 1 ℎ ℎ → 𝜆 = 𝑅𝐻 (𝑛 2 − 𝑛 2 ) ; Rydberg’s Constant (R) = or → 𝜆 = 𝑚𝑣 = ; m or me is the mass of object 𝑓 𝑖 √2𝑚𝑒 𝑞𝑒 𝑉 CONST 16 (for Canon F-789SGA), and nf and ni are or electron (CONST 03), v is the velocity, qe is the the final and initial energy levels, respectively charge of the electron (CONST 23) and V is the voltage *CONST for Canon F-789SGA IV.IV QUANTUM NUMBERS Quantum Mechanical Model → used to distinguish the probability of finding electrons in 3D space according to mathematical function Nodes → regions in space without electron density Quantum Numbers → address of electrons in an atom Quantum Number Definition Principal (n = 1, 2, 3 Main energy level or Atomic Spectrum of Hydrogen → result of … 7) the distance of excitation of the atoms and releasing the excess electrons from the nucleus energy by emitting light of various wavelengths (Hydrogen Spectrum Series) Azimuthal (ls = 0, lp = Energy subshells or 1, ld = 2, lf = 3 the shape of the Hydrogen Spectrum Series orbitals Final Energy Level Series Magnetic (ml = -l to Number of orbitals in (nf) +l) subshell or the 1 Lyman possible orientation of orbitals in space 2 Balmer Spin (ms = -½, Movement of electrons 3 Paschen counterclockwise, around its own axis 4 Brackett downward or ½, 5 Pfund clockwise, upward) 6 Humphreys Maximum number of orbitals in shell: 𝑛2 IV.II LIGHT AS A PARTICLE Maximum number of electrons in shell: 2𝑛2 Albert Einstein → concluded that when photons are shone at a metal surface, some electrons can be IV.V ELECTRON CONFIGURATION knocked off in the metal structure (Photoelectric Effect) Electron Configuration → arrangement of electrons in the orbitals of an atom Kinetic Energy (KE) → energy that will drive away electrons from the metal surface 1 → 𝐾𝐸 = 2 𝑚𝑣 2 ; m is the mass of the electron and v is the velocity of the electrons Work Function (𝝓) → minimum energy to remove electrons from the metal surface → 𝜙 = ℎ𝑓0 ; Planck’s Constant (h) = or CONST 06 (for Canon F-789SGA) and f0 is the minimum or threshold frequency Energy of Incident Light (ETotal) → 𝐸𝑇𝑜𝑡𝑎𝑙 = 𝜙 + 𝐾𝐸 Aufbau Principle → also called the building up principle; orbitals are filled with electron in increasing IV.III ELECTRON AS A WAVE energy Louis de Broglie → postulated that electrons can also behave as waves Hund’s Rule of Maximum Multiplicity → the most This type of interaction arises when stable arrangement of electrons in subshells is the there is induction of a dipole in the one with more parallel spins molecule. → in same energy orbital, upward electrons should 3. Dispersion be filled up first before the downward electrons These are also called London forces or Van der Waals forces and exist Pauli’s Exclusion Principle → no two electrons mostly in covalent molecules. can have the same set of four quantum numbers 4. Hydrogen bonds an interaction between the Hydrogen Ground State → lowest energy arrangement of and an electronegative element electrons in the orbital of an atom where there is no electron exchange leading to weak bonding between Excited State → allowed arrangements other than molecules the ground state Octet rule and Lewis structure Valence Electron → electrons with the highest principal quantum number Lewis Electron Dot Symbols 20th century Core Electrons → electrons closer to the nucleus Gilbert N. Lewis (1875–1946) used for predicting the number of bonds Valence Electron Configuration → all the number formed by most elements in their of electrons of the highest energy level compounds. Each Lewis dot symbol consists of the Isoelectronic → Species with the same electron chemical symbol for an element surrounded configuration by dots that represent its valence electrons. Notations of Electron Configuration Octet Rule 1. Condensed → the usual electron states that atoms tend to gain, lose, configurations or share electrons in order to achieve 2. Expanded → all electrons of every a stable configuration with eight subshells are expanded with two valence electrons. electrons each 3. Noble Gas → simplified condensed Octet Rule Exceptions: electron configuration using the previous nearest noble gas to the element Expanded octet - elements that can accommodate CHEMICAL BONDS more than eight electrons in their valence shell, such as elements from Chemical bond: the third period onwards. The driving force for chemical bonding is to Incomplete octet attain stable electronic configuration. - elements that have fewer than eight A chemical bond may be formed either by electrons in their valence shell, such electron transfer or electron sharing. as hydrogen and helium. Odd number octet Types of chemical bonds - elements that have an odd number of 1. Ionic Bond valence electrons, such as nitrogen. Metal - Nonmetal Therefore, all of these options 2. Covalent Bond deviate from the octet rule. Nonmetal - Nonmetal 3. Hydrogen Bond Molecular geometry Hydrogen - Electronegative atom - the three-dimensional arrangement of all the 4. Metallic Bond atoms in a given molecule. Metal - Metal - Note that the type of bonds – single, double, or triple – doesn’t influence the geometry of Types of Interactions: a molecule. 1. Electrostatic Interactions between charged molecules. 2. Induction  Butene Molecular = C4H8 Steric number - the number of domains attached to 4 Carbon atoms a central atom (atoms and lone pairs) 8 Hydrogen atoms Polarity Empirical = CH2 - in a bond is caused by electronegativity 4 C= 1C differences between the bonded atoms. 8H 2H - Electronegativity can be found in periodic table If the ratio of atoms in the molecular formula can’t be simplified any further, the empirical formula Polar (𝛥EN > 0.5) becomes the same as the molecular formula. Nonpolar (𝛥EN < 0.5) Percentage Composition - the percent by mass of The greater the difference in electronegativity, the each element in a compound greater the polarity. 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 Intermolecular Forces of Attraction (IMFA) %𝑚𝑎𝑠𝑠 = 𝑥 100 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑 Example: H2O Molecules Gas Liquid Solid Mass of compound = 18.02 IMFA Weak Strong Very Mass of H2 = 2.02 Strong Mass of O2 = 16 Gas molecules have weak IMFAs 2.02 % Hydrogen = 18.02 𝑥100 = 11.2% are highly compressible and have low 16 densities % Oxygen= 18.02 𝑥100 = 88.8% Liquid molecules have stronger IMFAs Chemical Reactions - involve the interaction of are slightly compressible and have high chemicals to form new substances densities represented by a chemical equation The left side of the equation represents the Solid molecules have very strong IMFAs reactants; the right side represents the are almost incompressible and have high products. densities Stoichiometric coefficients indicate the STOICHIOMETRY relative amounts of reactants and products in the reaction. Mole - a mole represents a quantity of a number, Compound states are indicated by symbols: the same way a “dozen” does. (l) for liquid, (s) for solid, (g) for gas, and (aq) 1 dozen = 12 for an aqueous solution. 1 mole = 6.022 𝑥 1023 = Avogadro’s Number Types of Chemical Reactions Molecular and Empirical Formula 1. Combination Reaction - two or more substances combine to form one product. Molecular Formula - how many atoms of General Formula: A + B → AB each element are in a compound Patterns for Combination Reactions Metal + Nonmetal → Binary Compound Empirical Formula - the simplest or most Nonmetal + Oxygen → Nonmetal Oxide reduced ratio of atoms in a compound Metal + Water → Metal Hydroxide (base) Nonmetal oxide + Water → Oxyacid (acid) Metal oxide + Nonmetal oxide → salt  2. Decomposition Reaction - a compound is They are not fully consumed because there decomposed to form two or more substances are more of them than needed to react with General Formula: AB △→ A + B the limiting reagent. Patterns for Decomposition Reactions Excess reactants are often left over after the Hydrates △→ salt + water reaction ends. IA Bicarbonates △→ Carbonates + H2O (g) + CO2 How to Determine the Limiting and Excess IIA Bicarbonates △→ Metal oxide + H2O (g) Reactant: + CO2 Carbonates △→ Metal oxide + CO2 1. Write the Balanced Equation: Ensure the Chlorates △→ Chlorine + Oxygen chemical equation for the reaction is Metal oxide △→ Metal + Oxygen balanced. Water △→ Hydrogen + Oxygen 2. Convert Quantities to Moles: Convert the masses or volumes of the reactants to 3. Displacement Reaction - more active metal moles. can displace a less active metal, while a less 3. Calculate the Mole Ratio: Use the active one can’t displace the more active. balanced equation to find the mole ratio General form: AY + B → BY + A; where A and B between the reactants. are metals based on its activity series 4. Compare the Mole Ratio to Determine the 4. Metathesis (Double Displacement Reaction) Limiting Reactant: - the positive ions exchange partners with the ○ Calculate how much of each negative ions to form two new compounds. reactant is required based on the General Form: AY + BX → BY + AX mole ratio. All neutralization reactions involving acids ○ Identify which reactant runs out first and bases are actually metathesis reactions. by comparing the actual amounts Any carbonate, either in the solid state or available to the required amounts. aqueous solution, reacts with acid to form This reactant is the limiting reactant. water, carbon dioxide gas, and salt. 5. Calculate the Amount of Product 5. Neutralization Reaction Formed: Use the amount of the limiting Acid + Base → Salt + Water reactant to determine the amount of product Metal oxide + acid → Salt + Water formed. Nonmetal oxide + Base → salt + water 6. Determine the Excess Reactant: Subtract Ammonia + Acid → Ammonium salt the amount of the excess reactant that actually reacts from the total amount Chemical Stoichiometry available to find how much of it is left over. Stoichiometry - used to describe the quantitative SOLUTION relationships between the reactants and products in a chemical reaction. A homogeneous mixture consisting of solute and solvent. Solute Solvent Present in small Does not change its amount phase in the formation Dissolved substance of solution Limiting Reactant Dissolving medium The reactant that is completely consumed Solubility factors during a chemical reaction. It determines the maximum amount of 1. Nature of Solute and Solvent product that can be formed. A solute can only be dissolved in a solvent when Once the limiting reagent is used up, the they are alike. A general rule is “like dissolves like” reaction stops, even if other reactants are still available. 2. Temperature For solid and liquid: solute increases when Excess Reactant temperature is increased. The reactant(s) that remain after the reaction For a gaseous: solute to a liquid solvent decreases has completed. as temperature increases. 3. Pressure The effect of pressure is only applicable for the 𝑚𝑜𝑙 𝑠𝑜𝑙𝑢𝑡𝑒 𝑀𝑜𝑙𝑎𝑙𝑖𝑡𝑦 (𝑚) = solubility of gasses in liquids. The higher the 𝑘𝑔 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 pressure of a gas, the more soluble it is. Normality Types of solution 𝑒𝑞𝑢𝑖𝑣𝑎𝑙𝑒𝑛𝑡𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 1. Unsaturated Solution - solvent can still dissolve 𝑁𝑜𝑟𝑚𝑎𝑙𝑖𝑡𝑦 (𝑁) = = 𝑛𝑀 𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 the solute Equivalents per Mole 2. Saturated Solution - if a solvent can’t no longer 𝑒𝑞 dissolve a given solute at a given temperature 𝑛( ) = 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝐻 + , 𝑂𝐻 − , 𝑜𝑟 𝑒 − 𝑚𝑜𝑙 3. Supersaturated Solution - if the solvent can’t dissolve the solute and need to be heated for it to be INTRODUCTION TO CHEMICAL EQUILIBRIUM dissolved Chemical Equilibrium → Chemical reactions tend Solubility rules in Water at 25 deg C to move towards a dynamic equilibrium in which both reactants and products are present but have no Insoluble further tendency to undergo net change. Soluble Compounds Compounds Equilibrium Law → Describes the relationship All nitrates, All carbonates, between the concentrations of reactants and bicarbonates, phosphates, products in a chemical reaction at equilibrium. It is chlorates chromates, and compounds and sulfides except expressed by the equilibrium constant (K) containing alkali metal that of alkali metal Equilibrium Constant (KC) → Ratio of the ions and ammonium ions and ammonium concentration of products to the concentration of the ion. ion. reactants, each raised to their respective All halides except that All hydroxides except stoichiometric coefficients. of Ag+, Hg2 2+ and Pb2+ that of alkali metal ions All sulfates except that and Ba++ Additionally: of Ag+, Ca++, Sr++, Ba++ and Pb++ Kc represents the equilibrium constant measured in moles per liter [𝐶]𝑐 [𝐷]𝑑 (concentration).𝐾𝑐 = [𝐴]𝑎[𝐵]𝑏 Concentrations of Solutions Weight/Weight Percent Kp represents the equilibrium constant calculated using the partial pressures of 𝑤 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 gasses. % = × 100 𝑤 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 [𝑝𝐶]𝑐 [𝑝𝐷]𝑑 𝐾𝑝 = [𝑝𝐴]𝑎 [𝑝𝐵]𝑏 Weight/Volume Percent 𝑤 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑔) Relationship Between Kc and Kp: % = × 100 𝐾𝑝 𝐾𝑐 = (𝑅𝑇)∆𝑛 𝑉 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (𝐿) Parts per Million (ppm) where ∆n = (total moles of gas on the product 𝜇𝑔 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑚𝑔 𝑠𝑜𝑙𝑢𝑡𝑒 side) - (total moles of gas on the reactant 𝑝𝑝𝑚 = = side) 𝑔 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 Parts per Billion (ppm) Characteristics of Equilibrium Constant 𝑛𝑔 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝜇𝑔 𝑠𝑜𝑙𝑢𝑡𝑒 1. Changes in concentration, pressure, 𝑝𝑝𝑏 = = 𝑔 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 temperature, or the presence of inert gasses can shift the equilibrium but do not change Molarity the equilibrium constant itself. 2. The equilibrium constant is related to the 𝑚𝑜𝑙 𝑠𝑜𝑙𝑢𝑡𝑒 𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 (𝑀) = standard free energy change (△G°) by the 𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 equation: △G° = -RT ln Kequ. Molality  3. The equilibrium constant for the reverse reaction is the reciprocal of the original 1 constant, i.e., Krev =. 𝐾𝑒𝑞𝑢 4. If the stoichiometry of the reaction changes, NUCLEAR CHEMISTRY AND NUCLEAR the equilibrium constant is raised to the ENERGY power corresponding to the change. 5. For a reaction A + B ⇌ C + D with constant NUCLEAR BINDING ENERGY AND NUCLEAR K, if the equation is multiplied by 4, the STABILITY equilibrium constant becomes K4. 6. In stepwise reactions leading to final Nuclear Binding Energy → the energy required to products, the net equilibrium constant is the separate a nucleus into neutrons and protons, which product of each individual step’s equilibrium are held with strong nuclear forces constant, K = K₁ × K₂ × K₃. → 𝐸 = 𝛥𝑚𝑐 2 ; Dm is the mass defect (difference of 7. For reactions with a common product, the the total mass of individual particles in the atom and equilibrium constant remains unchanged, but the actual mass of the atom) and c is the speed of higher concentrations of the common light product can decrease the concentration of other products. Radioactivity → emission of protons, neutrons, and electromagnetic waves from the nucleus of an FACTORS AFFECTING CHEMICAL unstable atom EQUILIBRIUM AND ITS RESPONSE Le Chatelier’s Principle → If a system at Radioactive Decay → process of losing energy equilibrium is subjected to a change in concentration through light emission of an unstable nucleus (C), pressure (P), or temperature (T), the system will adjust itself to counteract the disturbance and Nuclear Stability → determines whether the atom restore a new equilibrium. will undergo radioactive decay Summary of Le Chatelier’s Principle: Stable Unstable Even number of > 84 protons Value of protons and neutrons Change in Change Equilibrium Factor Equilibrium Reason Magic number of Neutron to proton to System Constant Position (Kc) protons and neutrons: ratio > 1 Extra 2, 8, 20, 28, 50 82, Shifts away concentration Increase from No change 126 needs to be substance used up C Need to produce more TYPES OF RADIOACTIVE DECAY Shifts of substance Decrease toward No change to make up for substance what was Types of Radioactive Decay removed 1. Alpha Decay →𝑎24 or 𝐻𝑒24 Shift For gas: → emission of alpha particles Pressure Increase towards side increase = No change → it has a positive charge, very low with fewer Volume penetrating power and very high ionizing moles of gas decrease P power. For gas: 0 0 Shift Pressure 2. Beta Decay → 𝛽−1 or 𝑒−1 towards side Decrease with more decrease = No change → emission of beta particles Volume moles of gas increase → it has a negative charge, intermediate Shifts away Extra heat / penetrating and ionizing power. Increase from heat / energy must Yes energy be used up 3. Gamma Emission → 𝛾00 More heat / → it has a no charge, very high T Shifts energy needs penetrating power and very low ionizing towards to be Decrease heat / produced to Yes power. 0 0 energy make up for 4. Positron Emission → 𝛽+1 or 𝑒+1 the loss 5. Electron Capture → 𝛾00 and 𝑒−10 Rates of both forward and Catalyst reverse and Inert - No Shift reactions are No change Gasses increased by the same amount KINETICS OF DECAY Half-life (t1/2) → also called as the decay time → uses the concept of first-order reaction to determine the time wherein the sample has lost half of its content. First-order Half-life Activity Reaction 𝑁 𝑙𝑛 2 = 𝑘𝑡1 𝑎 = 𝑘𝑁 𝑙𝑛 = −𝑘𝑡 2 𝑁0 → N is the remaining amount after time t, N0 id the initial amount, t is the time lapsed and k is the decay constant NUCLEAR REACTORS Nuclear Reactors → contains and control nuclear fission (splitting apart of atoms) that release energy in the form of heat Uranium-235 → a common nuclear fuel which is capable of producing nuclear reaction and it can readily undergo fission Parts of Nuclear Reactor 1. Core → contains the fuel elements and moderator; it is where the nuclear reaction occurs 2. Moderator → reduces speed of neutrons; usually light-water 3. Control Rods → control the rate of nuclear reaction by adsorbing excess neutrons; made up of Cadmium and Boron 4. Steam Generator → a heat exchanger that is used to produce the steam