CHEM 2004 Inorganic and Organic Chemistry PDF

UNIT 1: ATOMIC STRUCTURE AND CHEMICAL BONDING General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 MATTER PARTS OF AN ATOM anything that occupies space and has...

UNIT 1: ATOMIC STRUCTURE AND CHEMICAL BONDING General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 MATTER PARTS OF AN ATOM anything that occupies space and has mass Pure substances Mixtures fixed composition; a combination of two cannot be further or more pure purified substances Elements Homogeneous Matter cannot be subdivided by uniform composition chemical or physical means throughout Compounds Heterogeneous Matter elements united in fixed non-uniform composition positions HISTORY OF THE ATOM Greek philosophers proposed matter composition as early as 400-300 BC. Aristotle suggested matter is continuous and infinitely divided. Leucippus, Democritus, and Epicum proposed matter ATOMIC NUMBER (Z) composed of tiny, discrete, invisible, indivisible particles. represents the number of protons in the nucleus of an atom. In a These particles were called ATOMS, meaning indivisible or neutral atom, the number of protons is equal to the number undividing. of electrons. atoms comes from the word “atomos” which means “a” For a neutral atom: cannot be and “tomos” means divided. Atomic number (Z) = No. of Protons = No. of Electrons FUNDAMENTAL PARTICLES MASS NUMBER (A) the total number of protons and neutrons in the nucleus of an PROTON atom. The mass number is always a whole number. denoted by “p” or “p⁺” is a subatomic particle with a Mass number (A) = atomic number + no. of neutrons = (no. positive electric charge of +1 elementary charge. of protons/electrons) + no. of neutrons Named after the the Greek word meaning “first” and this name was given to the hydrogen nucleus. The number of protons in the nucleus is the defining property of an element, represented by the symbol Z. NEUTRON a subatomic particle, symbol n or n⁰, with no net electric charge. Mass slightly larger than proton. Comprises nuclei of atoms along with protons. ELECTRON Subatomic particle symbol e− or β− with -1 elementary charge. Essential in electricity, magnetism, chemistry, thermal conductivity. Participate in gravitational, electromagnetic, and weak interactions. E.R.A. 1 UNIT 1: ATOMIC STRUCTURE AND CHEMICAL BONDING General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 ISOTOPES CATEGORIES OF ELECTRONS are atoms of the same element having the same atomic 1. Inner (core) electrons fill all the lower energy levels of an number but different atomic masses. atom. Isotopes = same number of protons but different number of 2. Outer electrons are those in the highest energy level neutrons. (highest n value). They spend most of their time farthest from the nucleus. 3. Valence electrons are those involved in forming compounds. Among the main group elements, the valence electrons are the outer electrons. Among the transition elements, all the (n-1)d electrons are counted as valence electrons also, even though the elements Fe (Z = 26) through Zn (Z = 30) utilize only a few of their d electrons in bonding. ISOBARS THE 4 PRINCIPAL QUANTUM NUMBERS are atoms of different elements having the same mass Quantum numbers designate specific shells, subshells, number but different atomic numbers. orbitals, and spins of electrons. This means that they Examples: describe completely the characteristics of an electron in an atom. 1. Principal quantum number (n) - describes the size of orbitals and relative distance of the electrons from the nucleus. ISOTONES The higher the n-value, the bigger is the orbital size, and the are atoms of different elements having the same number of farther the electron is from the nucleus. neutrons but differ in atomic numbers and mass numbers. ex. n=1 first energy level Examples: n=2 2nd energy level n=3 3rd energy level n=4 4th energy level.... ORBITALS It is the space around the nucleus in which the electron is 2. Angular momentum quantum number (l) – this refers to found with a probability of 90%. the shape of the orbitals. l has values of 0, 1, 2, 3, 4... or n-1 The electron spends 90% of its time in that space. from the n value. Can accommodate a maximum of 2 electrons 3. Magnetic quantum number (mₗ)- refers to the orientation of SHELLS the orbitals in space around the nucleus. The number of electrons in an atom are arranged in shells or 'energy levels' around the nucleus. The arrangement of # of orbitals electrons determines chemical properties of an element. The electron shells are 1, 2, 3, 4, 5, 6, and 7; going from n=1 l=0 mₗ = 0 1 innermost shell outwards. Electrons in outer shells have higher average energy and travel farther from the nucleus n=2 l=1 mₗ = +1, 0, -1 3 than those in inner shells. Electrons that occupy the first electron shell are closer to the n=3 l=2 mₗ = +2, +1, 0, -1, -2 5 nucleus and have a lower energy than electrons in the n = 4 l = 3 mₗ = +3, +2, +1, 0, -1, -2, -3 7 second electron shell. SUBSHELL- Each shell is composed of one or more subshells, 4. Spin quantum number (mₛ) - describes the spin of which are themselves composed of atomic orbitals. It is a electrons in an orbital which is opposite direction to differentiate region of space within an electron shell that contains one electron from the other in the same orbital. It has only two electrons that have the same energy. possible values. mₛ = + 1/2 mₛ = - 1/2 E.R.A. 2 UNIT 1: ATOMIC STRUCTURE AND CHEMICAL BONDING General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 Degenerate orbitals - Electron orbitals that have the same ELECTRONIC CONFIGURATION energy levels. the distribution of electrons in the energy levels and sublevels of an atom and is represented by the model: Orbitals in the 2p sublevel are degenerate - in other words the 2px, 2py, and 2pz orbitals are equal in energy, as shown in the diagram. SHELLS Likewise, at a higher energy than 2p, the 3px, 3py, and 3pz orbitals are degenerate. And, at a still higher energy, the 3dxy, 3dxz, 3dyz, 3dx2 - y2, and 3dz2 are degenerate. The number of different states of equal energy is called the degree of degeneracy or just degeneracy. The degeneracy of p orbitals is 3; the degeneracy of d orbitals is 5; the degeneracy of f orbitals is 7. RULES IN WRITING ELECTRONIC CONFIGURATION 1. Aufbau (building–up) principle: Electrons first occupy the lowest energy orbitals available to them; only when the lower-energy orbitals are filled that they enter higher – energy orbitals. 2. Pauli’s exclusion principle: only two electrons having opposite spins can occupy an orbital. The third electron will eventually be repelled. 3. Hund’s rule: electrons distribute singly before pairing. E.R.A. 3 UNIT 1: ATOMIC STRUCTURE AND CHEMICAL BONDING General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 n = group no. Paramagnetism - is a property of a substance to be slightly BOX DIAGRAM attracted to a magnetic field. It is exhibited by substances that contain one or more unpaired electrons. Diamagnetism - is a property of a substance to be slightly repelled by a magnetic field. It is exhibited by substances whose electrons are all paired. STATES OF AN ATOM GROUND STATE the lowest energy state or most stable state of an atom, molecule, or ion. Arranges electrons around atom's nucleus with lower energy levels. Electrons in varying energy orbitals naturally fall towards the lowest energy state. EXCITED STATE State other than ground state of an atom or molecule. Higher energy than ground state. Excitation is an elevation above baseline energy state. BOND FORMATION NOBLE GAS CONFIGURATION Chemical bond - the electrostatic force which holds the these gases have complete e- configuration of ns²np⁶ except for atoms in a compound or molecule. This results from the gain He making them difficult to either gain or loss electrons. and loss of electrons, or from the sharing of electrons between atoms. Valence is the ability of an atom to form bonds. The valence of an element is determined by the number of electrons in the outermost level of the atom. E.R.A. 4 UNIT 1: ATOMIC STRUCTURE AND CHEMICAL BONDING General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 BOND FORMATION COVALENT bonding Forms when atoms share electron pairs to form covalent OCTATE RULE molecules. The octet rule states that atoms are most stable when they Atoms share valence electrons to create a noble gas have a full shell of electrons in the outside electron shell. configuration. First shell has two electrons in a single s subshell. Bond results from mutual attraction for shared electrons. Helium has a full shell, making it stable and inert. Takes place between non-metals groups 4-7. Other shells have an s and p subshell, giving them at least Classified into polar and non-polar covalent bonds. eight electrons. Polar Covalent Bond Atoms combine to revert to the noble gas configuration with Forms when atoms have significant differences in eight electrons in the outermost shell. electronegativity values. Shared electron pair is closer to one atom than to the other. One end of the bond has partial positive charge, the other has partial negative charge. Bond has partial ionic character. Example: HCl bond. Non-Polar Covalent Bond Results from equal or nearly equal electron sharing by bonded atoms. Forms from combination of same element atoms or closely related atoms. The properties of a substance can be explained in terms of the Examples include H₂ bonds, Cl₂ bonds, N₂ bonds, and nature of the bonds holding the atoms together. C-H bonds in CH₄. TWO GENERAL TYPES OF CHEMICAL BONDS IONIC or electrovalent bonding Forms when electrons transfer from one atom to another. Results from mutual attraction of oppositely charged ions. Example: NaCl formation where sodium loses one valence electron, forming Na¹⁺ ion, and chlorine gains one, forming Cl¹⁻ ion with argon electronic configuration. Ions are electrically charged atoms of an element or group. Cations are positively charged atoms formed when an atom loses an electron. Anions are negatively charged atoms produced when an A covalent bond where both electrons come from the same atom gains an electron. atom. Group IA and Group IIA elements are most likely to form Forms when two atoms with similar or low electronegativity cations in ionic compounds. difference react. Examples of monoatomic ions: Cl¹⁻, K¹⁺, Fe²⁺, Bi⁵⁺ Allows both atoms to obtain noble gas electronic configuration. Examples of polyatomic ions: NO₂¹⁻, CO₃²⁻, PO₄³⁻ Coordinate covalent bond is when two electrons in the bond Cation is smaller than their neutral parent atom. are only donated by a single atom. Anion is larger than their parent atom as electrons are added to the atom, hence increasing the electron cloud size. E.R.A. 5 UNIT 1: ATOMIC STRUCTURE AND CHEMICAL BONDING General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 LEWIS DOT STRUCTURE PROPERTIES OF COVALENT BOND 2. Bond strength - the shorter the bond, the greater is the bond Represents element's nucleus by symbol. strength. Valence electrons are represented by pairs of dots. ELECTRONEGATIVITY TO DETERMINE POLARITY OF A MOLECULAR BOND 3. Bond Polarity: depends on the electronegativity values of elements. When the electronegativity difference between the bonded polar bond- results from unequal sharing of electrons by the bonded atoms. atoms is: non polar bond- results from equal or almost equal sharing greater than 1.7 - bond is ionic of electrons by the bonded atoms. 0.5 to 1.7 - bond is polar less than 0.5 - bond is nonpolar HF and NaI are the exceptions. For HF, the electronegativity difference is 4.0- 2.1 = 1.9. The bond is expected to be ionic, but the bond is polar. NaI, the electronegativity difference is 1.6 but the bond in NaI 4. Bond angle: Covalent molecules are bonded to other atoms is ionic. by electron pairs. This repulsion of valence electrons causes covalent molecules to have distinctive shapes, known as the molecule's molecular geometry. IONIC VS. COVALENT BONDS Covalent Ionic Polarity low high A covalent bond forms between two non-metals with An ionic bond similar forms between a electronegativity, metal and a non- PHYSICAL PROPERTIES OF COVALENT COMPOUND where neither metal, where the Formation atom is strong Generally have low boiling and melting points. stronger non- enough to attract Do not conduct electricity due to lack of electron metals easily electrons, and for transporting particles. attract electrons stabilization, they Bond length is shorter with stronger bonds and more share electrons from the metal. electrons. from outer Triple bonds are stronger than double bonds, double bonds molecular orbit. are stronger than single bonds. PROPERTIES OF COVALENT BOND Shape definite shape no definite shape 1. Bond length - Coulomb's law states that the bond energy is inversely related to the bond length (r), and so factors which Melting Point low high influence a bond's strength influence its length. Sodium chloride Bond length increases going down the periodic table. Methane (CH4), (smaller atoms make smaller bonds) (NaCl), Examples Hydro Chloric H-F < H-Cl < H-Br < H-I Sulphuric Acid acid (HCl) Bond length decreases going across the periodic table. (H2SO4) (smaller atoms make smaller bonds) C-C > C-N > C-O One metal and For bonds between equivalent atoms, the greater bond Occurs Between Two non-metals one non-metal order the shorter the bond length C≡C < C=C < C–C E.R.A. 6 UNIT 1: ATOMIC STRUCTURE AND CHEMICAL BONDING General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 IONIC VS. COVALENT BONDS EXPRESSION OF ATTRACTIONS OF ATOMS FOR ELECTRONS Covalent Ionic IONIZATION POTENTIAL Ionization potential is the energy needed to remove an electron Ionic bonds, also from a neutral atom. Metals lose electrons more easily than nonmetals, while known as nonmetals lose electrons more difficult. electrovalent Ionization energy increases over time due to the Shielding bonds, are Effect, with energy decreasing as atom size increases. Covalent bonding formed through The energy is required to attract all electrons of an atom to the is a chemical nucleus, forming a cation. electrostatic bonding between Chemical bond formation relies on the energy required to attraction non-metallic remove an electron. between What is it? atoms, involving oppositely the sharing of charged ions in a electron pairs chemical and other compound, covalent bonds. primarily between metallic and non-metallic atoms. Boiling Point low high State at Room Temperature liquid, gaseous solid ELECTRON AFFINITY COVALENCY VS. ELECTROCOVALENCY Electron affinity refers to the energy released when an atom gains one electron. Electrovalency refers to the number of electrons an atom loses It increases from left to right across a period and from bottom or gains to form ionic bonds, while covalency is the maximum to top along a group. number of electrons an atom can share to achieve a stable High for nonmetals but low for metals. electronic configuration, such as in the case of carbon. In many cases, electron affinity is positive, indicating energy release when an electron attaches to an atom. If the electron has to start a new shell due to full orbitals, it remains far from the nucleus and strongly repelled by existing electrons, resulting in negative electron affinity. ELECTRONEGATIVITY Electronegativity measures an atom's ability to attract electrons. Increases up a group of elements. Increases right in a period of elements. F is the most electronegative element followed by O. Low ionization energies and electron affinities result in low electronegativities. E.R.A. 7 UNIT 2: SYMBOLS, FORMULA, AND CHEMICAL EQUATION General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 CHEMICAL FORMULA TYPES OF INORGANIC COMPOUNDS A chemical formula is a way of presenting information about the chemical proportions of atoms that constitute a B. SALTS particular chemical compound or molecule, using chemical 1. Binary salts: are metals with multiple valence states, with element symbols, numbers, and sometimes also other symbols, their names followed by their Roman numeral or a suffix (ous) or such as parentheses, dashes, brackets, commas and plus and "ic" depending on the element's valence state. minus signs. Examples: CuS copper (II) sulfide or cupric sulfide TYPES Cu₂S copper (I) sulfide or cuprous sulfide 1. EMPIRICAL FeCl₃ iron (III) chloride or ferric chloride a formula giving the proportions or ratios of the elements FeCl₂ iron (II) chloride or ferrous chloride present in a compound but not the actual numbers or 2. Ternary salts: are compounds with more than two elements, arrangement of atoms. often bonding with polyatomic cations and anions, following Ex. CH₂O, HCO₂ rules similar to binary compounds in formulating. 2. MOLECULAR Examples: list the actual total no. of atoms present in a compound. Na₂SO₄ – sodium sulfate Ex. C₆H₁₂O₆ , H₂C₂O₄ K₃Fe(CN)₆ – potassium ferricyanide AlFe(CN)₆ – aluminum ferricyanide 3. STRUCTURAL Al₄[Fe(CN)₆]₃ – aluminum ferrocyanide gives the graphic representation of how atoms are arranged (NH₄)₂Cr₂O₇ – ammonium dichromate in a molecule or compound. BaCrO₄ – barium chromate Ex. BINARY COVALENT COMPOUNDS These compounds consist of two nonmetals, each named with a numerical prefix (di, tri, tetra, penta) indicating the number of atoms in the molecule. Examples: CO carbon monoxide CO₂ carbon dioxide STEPS IN WRITING A CHEMICAL FORMULA PCl₃ phosphorus trichloride 1. Identify elements/ions forming a compound. PCl₅ phosphorus pentachloride 2. Write valencies of elements/ions as superscript. CCl₄ carbon tetrachloride 3. Interchange valencies and write as subscript. P₂O₅ diphosphorus pentoxide 4. Write radical in parenthesis or brackets before subscript for P₂O₃ diphosphorus trioxide ions with different valencies. BASES 5. Write simple ratios of valencies if applicable. compounds that contain OH anion in the formula. 6. Write final formula without charge sign. cation + hydroxide NOMENCLATURE OF IONIC AND COVALENT COMPOUNDS Examples: Ionic compounds are formed when metals and nonmetals Al(OH)3 – aluminum hydroxide combine, with the cation being the first element and the NaOH- sodium hydroxide nonmetal anion being the second. NH4OH- ammonium hydroxide Ba(OH)2- barium hydroxide cation (specifying the charge, if necessary) + anion (element Ca(OH)2- calcium hydroxide stem + -ide) OXIDES a chemical compound that contains at least one oxygen atom and one other element. Metal oxides thus typically contain an anion of oxygen in the oxidation state of -2. Examples: TYPES OF INORGANIC COMPOUNDS A. ACIDS hydro + ion + ic + acid for binary acids ex. HCl – hydrochloric acid HF- hydrofluoric acid ACID SALTS are formed when replaceable hydrogen is partially replaced - ion + ic/ous + acid for ternary acids by a metal, with acids with two, three, or more hydrogens ex. forming one salt, and neutral salts with hydrogen or HNO₃ – nitric acid dihydrogen prefixes. HNO₂ – nitrous acid Examples: H2SO₄- sulfuric acid H₃PO₄ – phosphoric acid H₂CrO₃- chromous acid H₂CrO₄- chromic acid use –ic for higher valence, -ous for lower valence E.R.A. 1 UNIT 2: SYMBOLS, FORMULA, AND CHEMICAL EQUATION General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 TYPES OF INORGANIC COMPOUNDS TYPES OF SIMPLE CHEMICAL REACTIONS HYDRATES SINGLE DISPLACEMENT a substance that contains water on its constituent elements. between positively charged: A⁺+BC → AC+B⁺ Examples: between negatively charged: D⁻ + BC → BD + C⁻ CHEMICAL REACTIONS REACTIVITY OF METALS Chemical reaction is a process by which one or more chemical substances are converted into one or more different substances. Chemical Equations- a symbolic representation of a chemical reaction in the form of symbols and formula. A chemical reaction is described by a chemical equation Chemical equation provides identities and quantities of involved substances. Reactants (starting compounds) on left, products (final compounds) on right. Reactions are separated by an arrow. REACTIVITY OF NON METALS TYPES OF SIMPLE CHEMICAL REACTIONS DIRECT COMBINATION OR SYNTHESIS A + B → AB “FCBOIS” DOUBLE DISPLACEMENT AB + CD → CB + AD ANALYSIS OR DECOMPOSITION AB → A+B usually involves heating or electrolysis for chemical separation. COMBUSTION REACTIONS combustion reactions, typically involves hydrocarbons reacting with oxygen to produce carbon dioxide and water. E.R.A. 2 UNIT 2: SYMBOLS, FORMULA, AND CHEMICAL EQUATION General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 ENERGY CHANGES IN REACTIONS ENDOTHERMIC REACTIONS reactions in which energy is absorbed as heat. If heat is absorbed from the surroundings to the system, the surroundings would feel cold.. Examples: EXOTHERMIC REACTIONS reactions in which energy is evolved as heat. If heat is released from the system to the surroundings, the surroundings would feel hot. In thermodynamics, the term exothermic process (exo- "outside") describes a process or reaction that releases energy from the system to its surroundings, usually in the form of heat, light or sound.. Examples: E.R.A. 3 UNIT 3: NUCLEAR CHEMISTRY General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 NUCLEAR CHEMISTRY HISTORY OF RADIATION NUCLEAR CHEMISTRY ERNEST RUTHERFORD (1899) Subfield of chemistry focusing on radioactivity and nuclear Identified and classified three types of radiation: alpha, beta, processes. and gamma. Study of natural and artificially induced nuclear reactions. Differentiated between these types based on their NUCLEAR REACTION penetrating power and other properties. A reaction in which changes occur in the nucleus of an atom (not ordinary chemical reactions). TYPES OF RADIATION NUCLIDE ALPHA PARTICLES An atom with a specific atomic number and a specific Alpha particles have very low penetrating power. mass number. Can be stopped by a piece of paper or a thin sheet of RADIOACTIVE aluminum foil. a substance like uranium that spontaneously give off Alpha particles are the most ionizing type of radiation, radiation. meaning they cause significant ionization in the materials RADIATION they pass through. the penetrating rays and particles emmited by a Composed of two protons and two neutrons, identical to radioactive source. helium nuclei. Represented by the helium nucleus symbol or α. RADIOACTIVITY Carry a charge of +2. the emission of particles caused by the spontaneous Have a mass of 4 amu. disintegration of atomic nuclei. The emission of particles (alpha, beta, gamma/neutron) can occur spontaneously due to the decay of certain nuclides or due to changes in their internal structure. BETA PARTICLES Beta particles have moderate penetrating power. Can pass through a sheet of paper but are typically stopped by heavy clothing or a thin sheet of metal. Beta particles are less ionizing than alpha particles but still ionize the materials they pass through. Beta particles are identical to electrons. Represented by the symbol e⁻ or. Carry a charge of -1. Have a mass of approximately 1/1837 amu (practically negligible compared to other particles). HISTORY OF RADIATION WILHELM CONRAD ROENTGEN (1895) Accidentally discovered X-rays, a form of radiation more penetrating than ultraviolet rays. GAMMA RAYS Called the radiation "X-ray" to denote its unknown nature. Gamma rays are the most penetrating form of radiation. Initially thought to be a new type of invisible light; later Can pass through the body and many materials, causing recognized as a novel, unrecorded phenomenon. cellular damage. Effective shielding typically requires dense materials such as lead or thick concrete. HENRI BECQUEREL (1896) Consist of high-energy electromagnetic radiation with no Found that uranium salts emitted radiation spontaneously, mass or charge. leading to the discovery of natural radioactivity. High-energy electromagnetic radiation. Determined that it was the uranium material itself that was Sometimes referred to as neutron radiation or emission, producing the radiation, distinct from Roentgen’s X-rays. though this can be misleading as gamma rays are not MARIE CURIE neutrons. Coined the term "radioactivity" and made significant Gamma rays are less ionizing compared to alpha and beta discoveries in the field. particles but still pose significant risks due to their Identified and isolated radioactive elements such as penetrating ability. thorium, polonium, and radium. Represented by the symbol Shared the 1903 Nobel Prize in Physics with Becquerel and Carry no charge (0). Pierre Curie. Have no mass (0 amu). Awarded the 1911 Nobel Prize in Chemistry for her work on radioactivity. The element Curium was named in honor of Marie and Pierre Curie. Died from leukemia, attributed to prolonged exposure to radiation. E.R.A. 1 UNIT 3: NUCLEAR CHEMISTRY General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 OTHER TYPES OF RADIATION DETECTION AND MEASUREMENT OF RADIATION INTENSITY POSITRON DECAY Intensity refers to the energy flux or the number of particles A proton in the nucleus is converted into a neutron, or photons emerging per unit time. releasing a positron (the antimatter counterpart of an Measurement Instruments: electron) and a neutrino. Geiger-Muller Counter: Detects radiation by ionizing The atomic number of the atom decreases by one, while the gas within a tube, causing a current pulse that is mass number remains unchanged. counted. Decay Type: Often referred to as beta plus (β⁺) decay. Proportional Counter: Measures radiation by detecting Emission: A positron (β⁺) and a neutrino are emitted from the ionization within a gas-filled chamber, where the amount nucleus. of ionization is proportional to the energy of the radiation. Scintillation Counters: Use materials that emit flashes of light (scintillate) when struck by radiation; the light is then converted into an electrical signal and counted. Units of Measurement: Curie (Ci): Measures radioactivity; 1 Ci = 3.7 x 10¹⁰ NEUTRON CAPTURE disintegrations per second (dps). An atomic nucleus captures one or more neutrons. Becquerel (Bq): Measures radioactivity; 1 Bq = 1 The nucleus becomes heavier as it absorbs the neutrons, disintegration per second (dps). forming a new, heavier isotope or a different element. ENERGY Involves the collision and merging of neutrons with a The energy of radiation affects its penetrating power. nucleus. Different types of radiation have varying energies and The atomic number remains unchanged, but the mass penetrating capabilities. number increases by the number of captured neutrons. Alpha Particles: Mass and Charge: Most massive and highly charged. Penetrating Power: Least penetrating; can be stopped by a piece of paper or thin metal. Beta Particles: Mass and Charge: Less massive and less charged than alpha particles. Penetrating Power: More penetrating than alpha particles; can pass through paper but is stopped by heavy clothing or thin metal. Gamma Rays: Mass and Charge: No mass and no charge. Penetrating Power: Most penetrating; requires dense ELECTRON CAPTURE materials like lead for effective shielding. An electron from the closest energy level falls into the nucleus, causing a proton to convert into a neutron. A neutrino is emitted in this process. Additional electrons may fall into the resulting vacant energy levels, leading to a cascade of electron transitions. Decreases by one (since a proton is lost and replaced by a neutron). Remains unchanged as the number of nucleons (protons and neutrons) stays the same. Emission: A neutrino is emitted, and the atomic number decreases while the mass number remains unchanged. Electron Transition: Electrons from outer shells may fill the vacancies created by the capture. RADIATION DOSIMETRY RELATED TO HUMAN HEALTH Units of Measurement: Roentgen (R): The amount of radiation delivered from a radiation source. Rad (SI:Gray (Gy)): The ratio between radiation absorbed by a tissue and that delivered to the tissue (1 rad = 0.01 Gy). Rem (SI:Sievert (Sv)): The ratio between the tissue damage caused by a rad of radiation and the type of radiation (1 rem = 0.01 Sv). E.R.A. 2 UNIT 3: NUCLEAR CHEMISTRY General Chemistry (Organic and Inorganic) | LEC CHEM 2004 | Stub 112 | BSMLS-1A | HAGUISAN | 1st SEM 2024 SYMPTOMS FROM ACUTE RADIATION EXPOSURE NUCLEAR WASTE DISPOSALS Low-Level Waste: Disposed in landfills. Geological Disposal: Deep burial of waste. Recovery and Reuse: Extraction of usable materials from waste. Storage: Spent fuels stored underwater or in dry casks. Transmutation: Converts harmful elements into less harmful ones. Space Disposal: Launching waste into space (practicality issues). NUCLEAR MEDICINE APPLICATIONS NUCLEAR MEDICINE Nuclear medicine is the use of radioactive isotopes as tools for both diagnosis and treatment of diseases. Medical Imaging: Uses isotopes to diagnose/detect diseases. The goal is to create a useful picture of a target tissue. Radiation Therapy: Targets and destroys pathological tissue. RADIOACTVE ISOTOPES USEFUL IN MEDICAL IMAGING E.R.A. 3