Written Report (Inorganic Chem & PhyChem) PDF

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

INORGANIC CHEMISTRY Chemical Change of Matter - occurs when a substance undergoes a chemical reaction and Matter - anything that takes up space and can be forms one or more new substances with different weighed. In other words, matter has volume and chemical 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 The development of atomic theory has evolved not conform to the shape of their container. over time as scientific understanding has Liquid - particles are more loosely packed than in progressed. solids, allowing them to flow around each other. This gives liquids an indefinite shape, enabling Atomic Models them to conform to the shape of their container. Solid Sphere (John Dalton, 1803) Gasses - particles have a lot of space between Atoms are invisible and atoms of them and high kinetic energy. Gasses have no elements are identical. Compounds definite shape or volume and will spread out are combinations of atoms. indefinitely if unconfined. Plasma - is likely the most common state of matter Plum Pudding / Thomson Model (J.J. Thomson, in the universe, as seen in stars like the sun. 1904) Plasma consists of highly charged particles with extremely high kinetic energy. Electrons are scattered throughout a spherical cloud of positively charged Bose-Einstein Condensate - occurs at extremely particles. low temperatures, near absolute zero. In this state, a group of atoms is cooled to such low temperatures that they occupy the same quantum Nuclear / Rutherford Model (Ernest Rutherford, state, behaving as a single quantum entity with 1911) 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.  Subatomic Discovered Electric Charge Mass Pierre and Marie Curie → Discovered the Particle elements radium and polonium Electron J.J. −19 −31 (e-) Thomson − 1. 602×10 9. 11×10 Henry Moseley → Determined that atomic number, Proton Ernest −19 −21 (p+) Rutherford 1. 602×10 1. 673×10 not atomic weight, is the correct basis for ordering Neutron James −21 elements in the Periodic Table. 0 1. 675×10 (no) Chadwick Johann Döbereiner → Identified the Law of Triads Atomic Symbol for elements with similar properties, where the atomic weight of the middle elements of the triad was 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. Number of protons = Atomic number Robert Boyle → Defined elements as pure Number of electrons = Number of protons substances that cannot be broken down into Number of neutrons = Mass number - simpler substances, laying the groundwork for Atomic number modern chemistry. Atomic Properties Antoine Lavoisier → Known as the "Father of Metallic Property - ability Modern Chemistry", he developed a modern of an atom to donate an system of naming chemical substances, identified electron and named oxygen and hydrogen, and established the Law of Conservation of Mass. Atomic Size - average distance between the PARTS OF THE PERIODIC TABLE 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 Ionization Energy – number 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 Electronegativity – ability of an atom to attract or to 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 Usually shiny, very dense, high Table structure. melting points. Can be stretched into thin wires Antoine Bequerel → First discovered radioactivity. (ductile), hammered into thin sheets  (malleable), and are excellent elements are used in electronics, lasers, and strong conductors of heat and electricity. magnets. In chemical reactions, they readily lose electrons to form positive ions. Actinides → Radioactive elements, many of which All are solid at room temperature are synthetic. Uranium and thorium are notable for except for mercury. their use in nuclear energy. 2. Nonmetals Significant Elements: 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 electricity. one electron or like a halogen by gaining In chemical reactions, they tend to one electron. gain electrons, forming negative Oxygen → A highly reactive non-metal, ions. essential for combustion and respiration. Forms oxides with most elements and is the 3. Metals (or Semimetals) most abundant element in the Earth's crust. Exhibit both properties of metals and Carbon → Can form stable bonds with itself nonmetals and other elements, creating a vast array of Semiconductors: Can conduct organic compounds. The backbone of electricity better than insulators but organic chemistry and life. not as well as conductors 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 Atomic Radius → Increases as you move down, for containing 4 out of seven diatomic elements (F2, Decreases as you move across (left to right). Cl2, Br2, I2) Atomic size increases down a group due to added Noble Gases (Group 8A/18) → Inert, colorless, electron shells, expanding the electron cloud. and odorless gasses with a complete valence Across a period, it decreases as the increased electron shell, making them extremely stable and number of protons in the nucleus exerts a stronger mostly non-reactive. They are used in lighting and pull on the electrons, reducing the atomic radius. as inert atmospheres in chemical reactions. Lanthanides → Rare earth metals, known for their magnetic and phosphorescent properties. These  most reactive nonmetal as it is not found in nature as a free element and reacts explosively with many substances. Electronegativity → Increases as you move up and across (left to right) TYPES AND NAMING OF COMPOUNDS Fluorine remains the most electronegative element. Nomenclature Prefixes Noble gasses, however, have an electronegativity rating of 0 due to their inherent stability, preventing Number Prefix Number Prefix 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 electrons between metals and nonmetals. Ionization Energy → Increases as you move up and across (left to right) Steps for Naming Ionic Compounds (Monatomic Ions): Depends on how tightly electrons are held by the nucleus; electrons closer to the nucleus are more Example: MgBr2 tightly bound and require more energy to be removed. 1. Name of Metal Ion: The ion’s name (Mg2+) is given in the periodic table as magnesium. Electron Affinity → Increases as you move up 2. Name the non-metal ion by ending the and across (left to right) element name with the suffix “ide”: The nonmetal ion is Br (Bromine). By changing Electron affinity generally becomes more negative the name to end with the suffix “ide”, it gives across a period as atoms more readily accept “bromide”. electrons to complete their valence shell, and 3. Write the name of the compound: becomes less negative down a group due to Magnesium bromide reduced nuclear attraction on added electrons. Steps for Naming Ionic Compounds (Transition Metallic Character → Increases as you move Metal + Nonmetal): down, Decreases as you move across (left to right) Example: FeCl3 Refers to how easily an element can lose an electron. Cesium is the most reactive metal, 1. Identify the Transition Metal and reacting explosively with water and igniting Nonmetal: Determine which transition metal spontaneously in air. Francium, located below (Fe) and nonmetal (Cl) are present in the cesium in the alkali metal group, is so rare that its compound. properties are largely unobserved. 2. Determine the Charges of the Ions: Transition metals can form cations with Nonmetallic Character → Increases as you move different charges. Find the charge of the up and across (left to right) transition metal ion (3+) from the formula of the compound and the charge of the Refers to the tendency of an element to gain nonmetal ion (1-). electrons in chemical reactions. Fluorine is the  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 the types of bonds (e.g., CH₄ is methane, C₂H₄ is name with the name of the nonmetal which ethene). is “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 to each other in the direction of propagation. is 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 a. Binary acids (containing hydrogen and one wavelength other element) are named with the "hydro-" −8 𝑚 → speed of light =3 × 10 prefix and "-ic" suffix (e.g., HCl is 𝑠 𝑚 hydrochloric acid). → speed of sound =343 𝑠 b. Oxyacids (containing hydrogen, oxygen, and another element) are named based on Max Planck → postulated that light and other the polyatomic ion; "-ate" becomes "-ic acid" electromagnetic waves can be quantized and were and "-ite" becomes "-ous acid" (e.g., H₂SO₄ emitted in terms of discrete packets of energy is sulfuric acid). called Quanta ℎ𝑣 E. Bases → Compounds that release OH⁻ ions in →𝐸 = ℎ𝑓 = λ ; Planck’s Constant (h) = or solution. Named by stating the metal cation CONST 06 (for Canon F-789SGA) followed by "hydroxide" (e.g., NaOH is sodium hydroxide). Rydberg’s Equation → use to quantify the wavelength when the atom transitions from the F. Salts → Ionic compounds formed from the neutralization reaction between an acid and a base. excited phase to ground phase after exposure to IV.III ELECTRON AS A WAVE heat/flame. → 1 = 𝑅𝐻( 1 2 − 1 2 ) ; Rydberg’s Constant (R) = Louis de Broglie → postulated that electrons can λ 𝑛𝑓 𝑛𝑖 also behave as waves or CONST 16 (for Canon F-789SGA), and nf and ni →λ= ℎ = ℎ ; m or me is the mass of object 𝑚𝑣 are the final and initial energy levels, respectively 2𝑚𝑒𝑞𝑒𝑉 or electron (CONST 03), v is the velocity, qe is the 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 Atomic Spectrum of Hydrogen → result of Principal (n = 1, 2, 3 Main energy level or excitation of the atoms and releasing the excess … 7) the distance of energy by emitting light of various wavelengths electrons from the (Hydrogen Spectrum Series) nucleus Azimuthal (ls = 0, lp = Energy subshells or Hydrogen Spectrum Series 1, ld = 2, lf = 3 the shape of the Final Energy Level Series orbitals (nf) Magnetic (ml = -l to Number of orbitals in 1 Lyman +l) subshell or the 2 Balmer possible orientation of 3 Paschen orbitals in space 4 Brackett Spin (ms = -½, Movement of 5 Pfund counterclockwise, electrons around its 6 Humphreys downward or ½, own axis clockwise, upward) IV.II LIGHT AS A PARTICLE 2 Maximum number of orbitals in shell: 𝑛 Albert Einstein → concluded that when photons are shone at a metal surface, some electrons can 2 be knocked off in the metal structure Maximum number of electrons in shell: 2𝑛 (Photoelectric Effect) IV.V ELECTRON CONFIGURATION Kinetic Energy (KE) → energy that will drive away electrons from the metal surface Electron Configuration → arrangement of 1 2 electrons in the orbitals of an atom → 𝐾𝐸 = 2 𝑚𝑣 ; 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 1. Electrostatic principle; orbitals are filled with electron in Interactions between charged increasing energy molecules. 2. Induction 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 2. Expanded → all electrons of every achieve a stable configuration with subshells are expanded with two eight 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 Chemical bond: from 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 1. Ionic Bond of valence electrons, such as Metal - Nonmetal nitrogen. Therefore, all of these 2. Covalent Bond options 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.  Butene Molecular = C4H8 4 Carbon atoms Steric number - the number of domains attached to 8 Hydrogen atoms a central atom (atoms and lone pairs) Empirical = CH2 Polarity 4 C= 1C - in a bond is caused by electronegativity 8H 2H differences between the bonded atoms. - Electronegativity can be found in periodic If the ratio of atoms in the molecular formula can’t table be simplified any further, the empirical formula becomes the same as the molecular formula. Polar (∆EN > 0.5) Nonpolar (∆EN < 0.5) Percentage Composition - the percent by mass of each element in a compound The greater the difference in electronegativity, the greater the polarity. %𝑚𝑎𝑠𝑠 = 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 𝑥 100 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑 Intermolecular Forces of Attraction (IMFA) Example: H2O Mass of compound = 18.02 Molecules Gas Liquid Solid Mass of H2 = 2.02 Mass of O2 = 16 IMFA Weak Strong Very Strong 2.02 % Hydrogen = 18.02 𝑥100 = 11. 2% 16 Gas molecules have weak IMFAs % Oxygen= 𝑥100 = 88. 8% 18.02 are highly compressible and have low densities Chemical Reactions - involve the interaction of chemicals to form new substances Liquid molecules have stronger IMFAs are slightly compressible and have high represented by a chemical equation densities The left side of the equation represents the reactants; the right side represents the Solid molecules have very strong IMFAs products. are almost incompressible and have high Stoichiometric coefficients indicate the densities relative amounts of reactants and products in the reaction. STOICHIOMETRY Compound states are indicated by symbols: (l) for liquid, (s) for solid, (g) for gas, and Mole - a mole represents a quantity of a number, (aq) for an aqueous solution. the same way a “dozen” does. 1 dozen = 12 Types of Chemical Reactions 23 1 mole = 6.022 𝑥 10 = Avogadro’s Number 1. Combination Reaction - two or more Molecular and Empirical Formula substances combine to form one product. General Formula: A + B → AB Molecular Formula - how many atoms of Patterns for Combination Reactions each element are in a compound Metal + Nonmetal → Binary Compound Nonmetal + Oxygen → Nonmetal Oxide Empirical Formula - the simplest or most Metal + Water → Metal Hydroxide (base) reduced ratio of atoms in a compound 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 can moles. displace a less active metal, while a less active 3. Calculate the Mole Ratio: Use the 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 by comparing the actual amounts reactions. available to the required amounts. Any carbonate, either in the solid state or This reactant is the limiting reactant. aqueous solution, reacts with acid to form 5. Calculate the Amount of Product water, carbon dioxide gas, and salt. Formed: Use the amount of the limiting 5. Neutralization Reaction reactant to determine the amount of product Acid + Base → Salt + Water formed. Metal oxide + acid → Salt + Water 6. Determine the Excess Reactant: Subtract Nonmetal oxide + Base → salt + water the amount of the excess reactant that Ammonia + Acid → Ammonium salt actually reacts from the total amount available to find how much of it is left over. Chemical Stoichiometry SOLUTION Stoichiometry - used to describe the quantitative relationships between the reactants and products in A homogeneous mixture consisting of solute and a chemical reaction. solvent. Solute Solvent Present in small Does not change its amount phase in the formation Dissolved substance of solution Dissolving medium Limiting Reactant Solubility factors The reactant that is completely consumed during a chemical reaction. 1. Nature of Solute and Solvent It determines the maximum amount of A solute can only be dissolved in a solvent when product that can be formed. they are alike. A general rule is “like dissolves like” Once the limiting reagent is used up, the reaction stops, even if other reactants are 2. Temperature still available. For solid and liquid: solute increases when temperature is increased. Excess Reactant For a gaseous: solute to a liquid solvent decreases The reactant(s) that remain after the as temperature increases. reaction has completed. 3. Pressure Normality The effect of pressure is only applicable for the 𝑒𝑞𝑢𝑖𝑣𝑎𝑙𝑒𝑛𝑡𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 solubility of gasses in liquids. The higher the 𝑁𝑜𝑟𝑚𝑎𝑙𝑖𝑡𝑦 (𝑁) = 𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 = 𝑛𝑀 pressure of a gas, the more soluble it is. Equivalents per Mole Types of solution 1. Unsaturated Solution - solvent can still dissolve 𝑛 ( ) = 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝐻 , 𝑂𝐻 , 𝑜𝑟 𝑒 𝑒𝑞 𝑚𝑜𝑙 + − − the solute 2. Saturated Solution - if a solvent can’t no longer INTRODUCTION TO CHEMICAL EQUILIBRIUM dissolve a given solute at a given temperature Chemical Equilibrium → Chemical reactions tend 3. Supersaturated Solution - if the solvent can’t to move towards a dynamic equilibrium in which dissolve the solute and need to be heated for it to both reactants and products are present but have be dissolved no further tendency to undergo net change. Equilibrium Law → Describes the relationship Solubility rules in Water at 25 deg C between the concentrations of reactants and products in a chemical reaction at equilibrium. It is Insoluble Soluble Compounds expressed by the equilibrium constant (K) Compounds All nitrates, All carbonates, Equilibrium Constant (KC) → Ratio of the bicarbonates, phosphates, concentration of products to the concentration of chlorates chromates, the reactants, each raised to their respective and compounds and sulfides except stoichiometric coefficients. containing alkali metal that of alkali metal ions and ammonium ions and ammonium Additionally: ion. ion. All halides except that All hydroxides except Kc represents the equilibrium constant of Ag+, Hg2 2+ and Pb2+ that of alkali metal ions measured in moles per liter (concentration). 𝑐 𝑑 All sulfates except that and Ba++ 𝐾𝑐 = [𝐶] [𝐷] 𝑎 𝑏 [𝐴] [𝐵] of Ag+, Ca++, Sr++, Ba++ and Pb++ Kp represents the equilibrium constant calculated using the partial pressures of gasses. Concentrations of Solutions 𝑐 𝑑 [𝑝𝐶] [𝑝𝐷] 𝐾𝑝 = 𝑎 𝑏 [𝑝𝐴] [𝑝𝐵] Weight/Weight Percent 𝑤 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 Relationship Between Kc and Kp: % 𝑤 = 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 ×100 𝐾𝑝 𝐾𝑐 = ∆𝑛 (𝑅𝑇) Weight/Volume Percent 𝑤 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑔) where ∆n = (total moles of gas on the % 𝑉 = 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (𝐿) ×100 product side) - (total moles of gas on the reactant side) Parts per Million (ppm) μ𝑔 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑚𝑔 𝑠𝑜𝑙𝑢𝑡𝑒 Characteristics of Equilibrium Constant 𝑝𝑝𝑚 = 𝑔 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 = 𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 1. Changes in concentration, pressure, Parts per Billion (ppm) temperature, or the presence of inert gasses can shift the equilibrium but do not 𝑛𝑔 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 μ𝑔 𝑠𝑜𝑙𝑢𝑡𝑒 𝑝𝑝𝑏 = 𝑔 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 = 𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 change the equilibrium constant itself. 2. The equilibrium constant is related to the Molarity standard free energy change (△G°) by the 𝑚𝑜𝑙 𝑠𝑜𝑙𝑢𝑡𝑒 equation: △G° = -RT ln Kequ. 𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 (𝑀) = 𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 3. The equilibrium constant for the reverse reaction is the reciprocal of the original Molality 1 constant, i.e., Krev = 𝐾𝑒𝑞𝑢. 𝑚𝑜𝑙 𝑠𝑜𝑙𝑢𝑡𝑒 𝑀𝑜𝑙𝑎𝑙𝑖𝑡𝑦 (𝑚) = 𝑘𝑔 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛  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, product of each individual step’s equilibrium which are held with strong nuclear forces constant, K = K₁ × K₂ × K₃. 2 → 𝐸 = ∆𝑚𝑐 ; Dm is the mass defect (difference of 7. For reactions with a common product, the the total mass of individual particles in the atom equilibrium constant remains unchanged, and the actual mass of the atom) and c is the but higher concentrations of the common speed of 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 through light emission of an unstable nucleus concentration (C), pressure (P), or temperature (T), the system will adjust itself to counteract the Nuclear Stability → determines whether the atom disturbance and restore a new equilibrium. will undergo radioactive decay Summary of Le Chatelier’s Principle: Stable Unstable Even number of > 84 protons Change in Value of protons and neutrons Change Equilibrium Factor to System Equilibrium Reason Constant Magic number of Neutron to proton 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 Need to C 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 4 4 1. Alpha Decay →𝑎2 or 𝐻𝑒2 Shift For gas: towards side Pressure → emission of alpha particles Increase with fewer increase = No change moles of Volume → it has a positive charge, very low gas decrease penetrating power and very high ionizing P Shift For gas: power. towards side Pressure 0 0 Decrease with more decrease = No change 2. Beta Decay → β−1 or 𝑒−1 moles of Volume gas increase → emission of beta particles Shifts away Extra heat / → it has a negative charge, intermediate Increase from heat / energy must Yes energy be used up penetrating and ionizing power. 0 More heat / 3. Gamma Emission → γ0 T Shifts energy needs Decrease towards to be Yes → it has a no charge, very high heat / produced to energy make up for penetrating power and very low ionizing the loss power. Rates of both 0 0 forward and 4. Positron Emission → β+1 or 𝑒+1 Catalyst reverse 0 0 and Inert - No Shift reactions are No change 5. Electron Capture → γ0 and 𝑒−1 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 = 𝑘𝑡 1 𝑎 = 𝑘𝑁 𝑙𝑛 𝑁0 =− 𝑘𝑡 2 → 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 PHYSICAL CHEMISTRY pressure exerted by the gasses on the container can result in an explosion. GAS LAWS 𝑃1 𝑃2 In equation: 𝑇1 = 𝑇2 , where 𝑃1 and 𝑃2 are the Ideal Gas Law - also known as the general gas initial and final pressure and 𝑇1 and 𝑇2 are the initial equation and final temperature. Equation of a hypothetical ideal gas which is a good approximation of the behavior of Combined Gas Law many gasses under many conditions. The combined gas law is the law which combines States that the product of the pressure and Charles’s law, Gay-Lussac’s law and Boyle’s law. the volume of one gram molecule of an ideal gas is equal to the product of the The simultaneous variations of volume with absolute temperature of the gas and the temperature and pressure complete the relationship universal gas constant. In equation: PV = among P, V, and T of any gas. nRT Formula: Ideal Gas Law Units Combined Gas Law −1 −1 When R = 8.314 𝐽 𝑚𝑜𝑙 𝐾 then: P is P = pressure 3 𝑃𝑉 T = temperature in kelvin in pascals (Pa), V is in 𝑚 , and T is in Kelvin 𝑘 = 𝑇 V = volume (K) −1 k = constant (units of energy When R = 0.08206 𝐿 𝑎𝑡𝑚 𝐾 then: P is divided by temperature) in atm, V is in L, and T is in Kelvin (K) When two substances are compared in two different conditions, the law can be stated as: Boyle’s Law - the pressure of a given mass of gas Pi = initial pressure is inversely proportional to its volume, provided the Vi = initial volume 𝑃𝑖𝑉𝑖 𝑃𝑓𝑉𝑓 temperature remains constant. = Ti = initial temperature 𝑇𝑖 𝑇𝑓 When a gas is compressed (volume is Pf = final pressure reduced), its pressure increases, and when Vf= final volume a gas is expanded (volume is increased), its Tf = final temperature pressure decreases. In equation: 𝑃1𝑉1 = 𝑃2𝑉2, where 𝑃1 is the KINETIC MOLECULAR THEORY initial pressure exerted by the gas, 𝑉1 is the initial volume occupied by the gas, 𝑃2 is the Attempts to explain the properties of gases and gas laws final pressure exerted by the gas, and 𝑉2 is the final volume occupied by the gas First proposed by Bernoulli (1738) and extended by Clausius, Maxwell, Boltzmann, Vander Waals Charles’ Law - the volume of an ideal gas is and Jeans directly proportional to the absolute temperature at 1. Gases are composed of minute discrete particles constant pressure. called molecules. When a gas is heated, its volume will expand given that the pressure is held 2. The molecules within a container are believed to steady, and when the gas is cooled, the be in ceaseless chaotic motion during which they volume will contract. collide with each other and with the walls of the 𝑉1 𝑉2 container. In equation: 𝑇1 = 𝑇2 , where 𝑉1 and 𝑉2 are 3. The molecular collisions must involve no energy the initial and final volumes of the gas and loss due to friction thus all molecular collisions are 𝑇1 and 𝑇2 are initial and final absolute perfectly elastic. temperatures, measured in Kelvin. 4. The absolute temperature is proportional to their average kinetic energy. Gay-Lusaac’s Law - the pressure exerted by a gas is proportional to the temperature of the gas when 5. The molecules do not exert any force of the mass is fixed, and the volume is constant. attraction or repulsion on one another except during When a pressurized aerosol can (such as a collisions. deodorant can or a spray-paint can) is heated, the resulting increase in the 6. The volume of the particles is negligible compared with the total volume of the gas. Molecular Velocity of Gases Reduced Volume 𝑉𝑟 = 𝑉 𝑉𝑐 Most Probable Velocity COMPRESSIBILITY FACTOR 2𝑅𝑇 𝑣𝑚𝑝 = 𝑀 Compressibility Factor (Z) → A dimensionless Average Velocity factor that describes how much a real gas deviates from ideal behavior. It indicates the ability of a gas 8𝑅𝑇 𝑣 = π𝑀 to be compressed at a certain condition. Root Mean Square Velocity 𝑃𝑉 𝑍= 𝑅𝑇 (P2-190 Perry’s 9th Ed.) 3𝑅𝑇 𝑉𝑎𝑐𝑡𝑢𝑎𝑙 𝑃𝑉𝑎𝑐𝑡𝑢𝑎𝑙 𝑣𝑟𝑚𝑠 = 𝑀 𝑍= 𝑉𝑖𝑑𝑒𝑎𝑙 = 𝑅𝑇 Collision Properties For ideal gasses, Z=1. Deviations from 1 indicate the extent to which a real gas deviates from ideal Number of collisions each gas molecule encounters behavior due to interactions and finite volume. per second: 2 2π𝑑 𝑣𝑁𝐴 Z VALUE DOMINATING FAVORS EXAMPLE 𝑠 = 𝑉 FORCES 1 (Ideal) - - - Mean Free Path >1 Repulsive Expansion H2 The mean distance traveled by a gas molecule between two successive collisions. 1 f>P Repulsive Expansion Critical Constant → Conditions wherein the