Atomic Structure and Atomic Properties PDF

# Atomic Structure ## Fundamental Particles - Electron (e°) - Proton (p+) - Neutron (n°) ## Other Particles - Positron - Meson (+, 0, -) - Neutrino - Pie Meson (0, +, -) ## Cathode Rays - Discovered by J.J. Thomson - Produced in a discharge tube with high voltage and low pressure - Invisible rays emitted from the cathode towards the anode, creating a luminous path - Negatively charged particles (electrons) - Deflected by electric and magnetic fields - Produce fluorescence and phosphorescence in certain materials ## Anode rays / Canal rays / Positive rays - Produced in a discharge tube with high voltage and low pressure - Emitted from the anode towards the cathode - Positively charged particles - Composed of ions of the gas inside the discharge tube - Dependence on the nature of the gas in the discharge tube ## X-rays - Discovered by Wilhelm Roentgen - Produced when cathode rays strike a metal target - Highly penetrating electromagnetic radiation - Have a high frequency (10¹⁹ Hz) - Have a high penetrating power ## Radioactivity - Discovered by Henry Becquerel - The spontaneous emission of α, β, γ radiations by unstable nuclei of certain atoms - These radiations cause changes in the atomic nucleus and the atom becomes more stable ## α-Particle - (He+2)4 - A helium nucleus with a charge of +2 - Atomic mass = 4 amu - Contains 2 protons and 2 neutrons ## β-Particle - Similar to an electron - Involved in beta decay where a neutron decays into a proton and an electron - If a neutron decays into a proton and an electron (β-decay) it is called a β particle. ## γ-Particle - A form of electromagnetic radiation - Released during radioactive decay - Essentially high-energy photons ## Fundamental Particles ### Electron - Discovered by J.J. Thomson - Named by Stony - Charge = -1.6 x 10⁻¹⁹ C - Mass = 9.1 x 10⁻³¹ kg ### Proton - Discovered by Goldstein - Named by Rutherford - Charge = +1.6 x 10⁻¹⁹ C - Mass = 1.672 x 10⁻²⁷ kg ## Atomic Models ### Thomson’s Model - Proposed by J.J. Thomson in 1904 - Describes the atom as a positively charged sphere with negatively charged electrons embedded in it - The electrons were considered to be distributed in a uniform manner throughout the sphere - Known as the “plum pudding model” - It fails to explain the scattering of alpha particles by a gold foil, - It does not explain the existence of isotopes, - It does not explain the line spectrum of hydrogen. ### Rutherford’s Model - Proposed by Ernest Rutherford in 1911 - Based on the results of the alpha particle scattering experiment by a thin gold foil. - The atom consists of a small, dense, positively charged nucleus at its center - Negatively charged electrons revolve around the nucleus in circular orbits - It explains the alpha particles scattering experiment results. - Fails to explain the stability of an atom. - It fails to explain the line spectrum of hydrogen. ### Bohr's Model - Proposed by Niels Bohr in 1913 - A modification of Rutherford’s model, based on Planck’s Quantum Theory - Electrons revolve around the nucleus in stationary circular orbits without radiating energy - The energy levels are quantized and are represented by principal quantum numbers (n) which can be any positive integer, - The electrons can jump from one energy level to another by absorbing or emitting energy - The frequency of the emitted radiation is given by, $ν = \frac{E_2 - E_1}{h}$ Where, * E₂ is the higher energy level, * E₁ is the lower energy level, * h is Planck's constant - Bohr's model explains the line spectrum and the stability of atoms. - Bohr’s model failed to explain the spectra of atoms having more than 1 electron. - It fails to explain the Zeeman effect and Stark effect. ### Bohr-Sommerfeld Model - An extension of Bohr’s model - Proposed by Arnold Sommerfeld in 1916 - Electrons can revolve around the nucleus in elliptical orbits - Added the concept of an azimuthal quantum number (l) to describe the shape of the orbit ### Wave Mechanical Model - Developed by Erwin Shrodinger in 1926 - Describes electrons as wave-like particles, with dual nature - Based on the de Broglie’s hypothesis stating that every moving particle exhibits wave-like properties. - The electrons are localized in space and are characterized by a set of wave functions, called orbitals. - Describes the probability of finding an electron at a given point in space. ## Quantum Numbers - Quantum numbers are used to describe the properties of an electron in an atom. - Each electron in an atom is defined by a unique set of quantum numbers. - There are four types of quantum numbers: - Principal quantum number (n) - Azimuthal quantum number (l) - Magnetic quantum number (ml) - Spin quantum number (ms) ### Principal Quantum Number (n) - Describes the electron's energy level or shell - It is the highest energy level of the electron, and determines the size of the shell. - n=1, 2, 3, 4, 5 ...∞ indicating the first, second, third shells, and so on. - Higher the value of n, higher the energy level. - For a particular shell, n = ∞ indicates ionization of the atom. ### Azimuthal Quantum Number (l) - Describes the shape of the electron's orbital and its angular momentum. - It represents the number of subshells in a given shell. - For a particular shell, the possible values of l are 0, 1, 2, …., n-1 - Each value of l corresponds to a different orbital shape: * l = 0, s-orbital (spherical) * l = 1, p-orbital (dumbbell) * l = 2, d-orbital (clover-leaf) * l = 3, f-orbital (complex) ### Magnetic Quantum Number (ml) - Describes the orientation of the orbital in space relative to an external magnetic field. - The number of orbitals in a subshell is determined by the magnetic quantum number. - For a particular value of l, the values of ml range from -l through 0 to +l, including 0, thus, 2l+1 orbitals are possible for a particular value of l. - For example: * l = 0, ml = 0, s-orbital * l = 1, ml = -1, 0, +1, 3 p-orbitals * l = 2, ml = -2, -1, 0, +1, +2, 5 d-orbitals ### Spin Quantum Number (ms) - Describes the intrinsic angular momentum of an electron. - This angular momentum is called spin angular momentum and is associated with the spin of the electron. - Electrons behave as though they are spinning, creating a magnetic dipole moment. - The spin quantum number is used to describe the direction of this magnetic dipole moment. - It can only have two values: +1/2 or -1/2, representing spin up and spin down. ## Aufbau principle - States that an electron can enter a subshell only if all the subshells with lower energy are completely filled. - The energy level of a subshell is determined by the (n+l) rule, where n is the principal quantum number, and l is the azimuthal quantum number. - The subshells are filled in the increasing order of the (n+l) value. For subshells having the same (n+l) value, the subshell with lower n value is filled first. ## Hund’s Rule - States that electrons will individually occupy each orbital within a subshell before doubling up in any one orbital. - The electrons will all have the same spin until all orbitals are half-filled. - Provides the lowest energy configuration with greater stability. ## Pauli’s Exclusion Principle - States that no two electrons in an atom can have the same set of four quantum numbers. - This implies that each orbital can hold a maximum of two electrons, and these electrons must have opposite spins. ## Electronic Configuration - Refers to the arrangement of electrons in different energy levels and subshells within an atom. - Provides information about the chemical properties of an atom, including its reactivity. ## Periodic Table - A tabular arrangement of chemical elements organized by their atomic numbers and recurring chemical properties. - Elements with similar chemical properties are grouped together into columns known as groups or families. - The rows are known as periods. ### Development of the Periodic Table - **Levizioer** first classified elements into metals and non-metals - **Proust’s** hypothesis that all elements are made up of hydrogen, which he considered the simplest element. - **Dobereiner’s **Triads grouped elements into sets of three, with the middle element being the average in terms of atomic weight and chemical properties. - **Newlands’ **Law of Octaves noted that every eighth property was similar to the first for elements ordered by atomic weight. - **Mendeleev**’s Periodic Law states that the properties of elements are periodic functions of their atomic masses. Mendeleev’s table arranged elements according to increasing atomic mass, leaving gaps for undiscovered elements. It was organized into rows (periods) and columns (groups) based on similar chemical properties. - **Moseley’s** contribution was the use of X-rays to determine the actual atomic numbers of elements. Moseley’s work led to the modern periodic table. - **Bohr-Bury’s **classification based on the electronic configuration of elements, with the elements grouped together based on the similarities in the electronic configurations of their outermost shell. ### Classification of elements in the Periodic Table - **S-block elements**: Elements with the outermost electron in the s-orbital. They are soft metals and have low ionization energies. They are highly reactive. They are found in Group 1 and Group 2 of the periodic table. - **P-block elements**: Elements with the outermost electron in the p-orbital. The outermost electron in the p-orbital. They are located in groups 13 to 18 of the periodic table. This group is known as the p-block elements. - **d-block elements**: Elements with the outermost electron in the d-orbital. They are located in Groups 3 to 12 of the periodic table. These elements are known as transition metals, including: - Scandium, Titanium, Vanadium, Chromium, Manganese, Iron, Cobalt, Nickel, Copper, Zinc - Yttrium, Zirconium, Niobium, Molybdenum, Technetium, Ruthenium, Rhodium, Palladium, Silver, Cadmium - Hafnium, Tantalum, Tungsten, Rhenium, Osmium, Iridium, Platinum, Gold, Mercury - Rutherfordium, Dubnium, Seaborgium, Bohrium, Hassium, Meitnerium, Darmstadtium, Roentgenium, Copernicium, Nihonium, Flerovium, Moscovium, Livermorium. - **f-block elements**: Elements with the outermost electron in the f-orbital. They are located in the two rows below the main body of the periodic table. They are known as inner transition metals. - These include: - Lanthanides: Cerium, Praseodymium, Neodymium, Promethium, Samarium, Europium Gadolinium, Terbium, Dysprosium, Holmium, Erbium, Thulium, Ytterbium, Lutetium - Actinides: Actinium, Thorium, Protactinium. Uranium, Neptunium, Plutonium, Americium, Curium - Berkeium, Californium, Einsteinium. Fermium, Mendelevium, Nobelium, Lawrencium, Rutherfordium, Dubnium, Seaborgium, Bohrium, Hassium, Meitnerium, Darmstadtium, Roentgenium, Copernicium, Nihonium, Flerovium, Moscovium, Livermorium. ## Atomic Size - The size of an atom is defined as the distance between the nucleus and the outermost shell containing electrons. - Atomic size can be measured using the following techniques: - Covalent radius: Half the distance between the nuclei of two identical atoms linked by a single covalent bond. - Metallic radius: Half the distance between the nuclei of two adjacent metal atoms in the crystal lattice - Van der Waals radius: Half the distance between the nuclei of two adjacent atoms in the solid-state. ## Factors affecting atomic size - **Number of electron shells**: The greater the number of electron shells, the larger is the atomic size. - **Effective nuclear charge (Zeff)**: In a neutral atom, the electrons are attracted towards the nucleus. The attraction is proportional to Zeff where Z is the number of protons in the atom. As the number of atomic shells increases, the shielding effect from the inner shells increases, and therefore, the effective nuclear charge decreases. As a result, the outer electrons experience a weaker attraction towards the nucleus and are further away from the nucleus. - **Shielding effect**: Core electrons shield the valence electrons from the full attraction of the nucleus. - **Nuclear charge**: The higher the nuclear charge, the stronger the attraction between the nucleus and the electrons. This results in a smaller atomic size. ## Periodic Trends in Atomic Size - **Across a Period**: The atomic size decreases from left to right. - **Down a group**: The atomic size increases from top to bottom. ## Ionic Size - The ionic size is the radius of an ion, which is the distance between the nucleus and the outermost electron shell of the ion. - Cations are smaller than their parent atoms because they have lost electrons and the remaining electrons are pulled closer to the nucleus. - Anions are larger than their parent atoms because they have gained electrons and the increased number of electrons results in greater electron-electron repulsion, pushing the electron cloud further away from the nucleus. ## Ionization energy (IE) - The energy required to remove the most loosely bound electron from a gaseous atom, to form a gaseous cation is called ionization energy. - The first ionization energy is the energy required to remove one electron from a neutral atom. - The second ionization energy is the energy required to remove the second electron from a singly charged ion. - The third ionization energy is the energy required to remove the third electron from a doubly charged ion. ## Factors Affecting Ionization Energy - **Nuclear charge**: The greater the nuclear charge, the stronger the attraction between the nucleus and the electrons and therefore, - **Atomic size**: As the atomic size increase, the ionization energy decreases. - **Shielding effect**: The more core electrons there are, the weaker the attraction between the nucleus and the valence electrons, and therefore, the lower the ionization energy. - **Penetration effect**: The greater the penetration of the outer electron shells into the inner shells, the higher the ionization energy. - **Electron configuration**: Electrons in half-filled or filled orbitals are more stable than electrons in partially filled orbitals. Therefore, elements with half-filled or filled orbitals have slightly higher ionization energies. ## Periodic Trends of Ionization Energies - **Across a Period**: Ionization energy increases across a period. - **Down a Group**: Ionization energy decreases down a group. ## Electron Affinity - The change in energy when an electron is added to a neutral gaseous atom is called electron affinity. - A negative electron affinity value indicates that energy is released when an electron is added to the atom, indicating stability. - A positive electron affinity value indicates that energy is required to add an electron to the atom. ## Factors Affecting Electron Affinity - ** Nuclear charge**: The stronger the nuclear charge, the greater the attraction for an extra electron. - **Atomic size**: As the atomic size increases, electron affinity decreases because the added electron is further away from the nucleus. - **Shielding effect**: As the number of core electrons increases, the shielding effect increases, and therefore, electron affinity decreases. - **Electronic configuration**: Atoms with half-filled or completely filled orbitals have lower electron affinities than atoms with partially filled orbitals. ## Periodic Trend of Electron Affinity - **Across a period**: Electron affinity generally increases across a period from left to right. - **Down a group**: Electron affinity generally decreases with increasing atomic number. - Electron affinity does not follow a regular trend down a group. ## Electronegativity - The tendency of an atom in a molecule to attract the shared pair of bonding electrons is known as electronegativity. - It is a measure of the relative attraction of an atom for electrons in a chemical bond. - A higher electronegativity indicates a stronger attraction for electrons. - Elements with high electronegativities are typically nonmetals, while elements with low electronegativities are typically metals. ## Factors Affecting Electronegativity - **Nuclear charge**: Higher the nuclear charge, stronger the attraction for the electron and therefore, higher the electronegativity. - **Atomic size**: As atomic size increases, electrons are farther away from the nucleus and thus lesser attraction, and therefore, lower electronegativity. - **Shielding effect**: As the number of core electrons increases, the shielding effect increases, and therefore, electronegativity decreases. - **Electron configuration**: Atoms with half-filled and completely filled orbitals have slightly higher electronegativities. ## Periodic Trend of Electronegativity - **Across a period**: Electronegativity increases across a period from left to right - **Down a group**: Electronegativity decreases down a group. ## Chemical Bonding - A chemical bond is a force of attraction that holds atoms together in a molecule or compound. - Chemical bonding involves the sharing or transferring of electrons among atoms. ### Types of Chemical Bonds - **Ionic bonds**: Involve the electrostatic attraction between oppositely charged ions, formed by the transfer of electrons from a metal to a nonmetal. - **Covalent bonds**: Formed by the mutual sharing of electrons by two nonmetal atoms to achieve a stable electronic configuration. - **Metallic bonds**: Occur between the atoms of metals. The electrons in metallic bonds are delocalized, resulting in a cloud of electrons shared by many positively charged metal ions, forming a strong bond. - **Hydrogen bonds**: A special type of dipole-dipole interaction that is important in biological systems. They are formed between hydrogen atoms and an electronegative atom (oxygen, nitrogen, or fluorine) that is bonded to another electronegative atom. ## Hybridization - Occurs when atomic orbitals mix to form new hybrid orbitals that have different shapes and energies from the original atomic orbitals. - Hybridization helps explain the shapes of molecules and the directions of their bonds. ### Types of Hybridization - **sp³ hybridization**: occurs when one s orbital and three p orbitals mix to form four equivalent sp³ hybrid orbitals. These hybrid orbitals are oriented tetrahedrally, resulting in bond angles of 109.5°. Examples: CH₄, NH₃, H₂O. - **sp² hybridization**: occurs when one s orbital and two p orbitals mix to form three equivalent sp² hybrid orbitals. The orbitals lie in a plane and are oriented at 120° to each other, resulting in a trigonal planar geometry. Examples: BF₃, CO₂. - **sp hybridization**: occurs when one s orbital and one p orbital mix to form two equivalent sp hybrid orbitals. These orbitals are linearly arranged and are oriented at 180° to each other. Examples: BeCl₂, C₂H₂. ## VSEPR Theory - The Valence Shell Electron Pair Repulsion Theory (VSEPR) predicts the shapes of molecules based on the repulsion between valence electron pairs. - The repulsion between the electron pairs is minimum when they are as far apart as possible. - The following are the key concepts in VSEPR theory: - An electron pair is considered to be a region of high electron density. - The electron pairs repel each other. - The greater the electron density, the greater the repulsion. - The preferred shape of a molecule is the one that minimizes the repulsion between the electron pairs around the central atom. ## Dipole Moment - A measure of the polarity of a molecule. It is a vector quantity, indicating the direction and magnitude of the molecular dipole moment. - A dipole moment is created when there is a separation of charge within a molecule, resulting in a positive and a negative end. - A molecule is considered polar if it has a net dipole moment. ## Intermolecular Forces - These forces are responsible for holding molecules together in liquids and solids. - These forces arise from the interactions between the partial charges of molecules that have a dipole moment. - The strength of the intermolecular forces increases with increasing electronegativity difference between the atoms. ### Types of Intermolecular Forces - **Dipole-Dipole Forces**: occur between polar molecules, where the positive end of one molecule is attracted to the negative end of another molecule. - **Hydrogen Bonding**: is a special type of dipole-dipole interaction that occurs between a hydrogen atom bonded to a highly electronegative atom (oxygen, nitrogen, or fluorine) and a lone pair of electrons on an adjacent oxygen. - **London Dispersion Forces**: occur between all molecules, both polar and nonpolar. They are very weak interactions that arise from temporary fluctuations in electron distribution in molecules. ## Bond Energies - The amount of energy required to break a particular bond in a molecule is called bond energy. - The bond energy is a measure of the strength of a chemical bond. - The higher the bond energy, the stronger the bond. ## Bond Length - The distance between the nuclei of two atoms held together by a chemical bond is called bond length. - The bond length is affected by several factors including: - The type of atoms involved in the bond - The multiplicity of the bond (single, double, or triple) - The size of the atoms involved in the bond. ## Molecular Geometry - The three-dimensional arrangement of atoms in a molecule or ion is called molecular geometry. - It is determined by the arrangement of electron pairs around the central atom. - The molecular geometry affects the physical and chemical properties of the molecule or ion. ## Hybridisation - The mixing of atomic orbitals to form a new set of hybrid orbitals with equal energy is called hybridisation. - Hybrid orbitals help in explaining the geometry and the bonding characteristics of molecules ## VSEPR Theory - Valence Shell Electron Pair Repulsion Theory (VSEPR) is a model used to predict the shapes of molecules based on the repulsion between electron pairs in the valence shell of the central atom. - The theory states that the electron pairs arrange themselves around the central atom to minimize the repulsion between them. ## Bond Polarity and Dipole Moments - A polar bond is a covalent bond between two atoms that have different electronegativities. - The more electronegative atom attracts the shared electrons more strongly, resulting in a partial negative charge on that atom and a partial positive charge on the other atom, creating a dipole moment. - A molecule with a net dipole moment is considered to be polar. - A molecule with zero dipole moment is considered nonpolar. ## Intermolecular Forces - Intermolecular forces are attractive forces between molecules. - They are weaker than the intramolecular forces that hold atoms together in a molecule. - The type and strength of the intermolecular forces influence the physical properties of a substance, such as melting point, boiling point, and vapor pressure. ## Types of Intermolecular Forces - **Dipole-dipole forces**: They are attractive forces between polar molecules having permanent dipoles. - **Hydrogen bonding**: A special type of dipole-dipole interaction that involves a hydrogen atom bonded to a highly electronegative atom like oxygen, nitrogen, or fluorine. - **London dispersion forces**: These are temporary, weak intermolecular forces that arise from the instantaneous imbalance of electron distribution around an atom. They are present in all substances, both polar and nonpolar. ## Solubility - The ability of a substance (solute) to dissolve in another substance (solvent) is called solubility. - The solubility of a substance depends on several factors, including: - The nature of the solute and solvent - Temperature - Pressure ## Chemical Reactions - Chemical reactions involve the rearrangement of atoms and molecules. - The study of the rate and extent of chemical reactions is called chemical kinetics. ## Types of Chemical Reactions **Based on energy changes** - **Exothermic reactions**: Release energy to the surroundings. - **Endothermic reactions**: Absorb energy from the surroundings. **Based on the number of reactants and products** - **Decomposition reactions**: One reactant breaks down into 2 or more products. - **Combination reactions**: Two or more reactants combine to form a single product. - **Single displacement reactions**: One element replaces another element in a compound. - **Double displacement reactions**: Two compounds exchange ions to form two new compounds. ## Factors Affecting Reaction Rate - **Concentration of reactants**: Increasing concentration of reactants generally increases the reaction rate. - **Temperature**: Increasing the temperature usually increases the reaction rate because molecules have more energy and move faster. - **Surface area**: Increasing the surface area of the reactants increases the reaction rate. - **Catalyst**: A catalyst speeds up the rate of a reaction without itself being consumed in the process. ## Chemical Equilibrium - The state where the rates of the forward and reverse reactions are equal, and the concentrations of reactants and products remain constant is called chemical equilibrium. - It is a dynamic process that occurs when a reversible reaction reaches a state where the rate of the forward reaction equals the rate of the reverse reaction. ## Equilibrium Constant (K) - K is a constant that indicates the ratio of product concentrations to reactant concentrations at equilibrium. - It is a measure of the extent to which a reaction proceeds to completion. ## Le Chatelier’s Principle - States that if a change in conditions is applied to a system at equilibrium, the system will shift in a direction that relieves the stress. - The stress can be a change in temperature, pressure, or concentration. ## Acids, Bases, and Salts - **Acids**: Substances that donate hydrogen ions (H⁺) when dissolved in water. - **They have a pH value less than 7.** - **Examples**: HCl, HNO₃, H₂SO₄ - **Bases**: Substances that accept hydrogen ions (H⁺) or donate hydroxide ions (OH⁻) when dissolved in water. - **They have a pH value greater than 7.** - **Examples**: NaOH, KOH, Ca(OH)₂. - **Salts**: Ionic compounds formed by the reaction of an acid and a base. - They are neutral and have a pH value of 7. - **Examples**: NaCl, KBr, CaCO₃. ## Acid-Base Theories ### Arrhenius Theory - **Acids** are substances that produce hydrogen ions (H⁺) in solution. - **Bases** are substances that produce hydroxide ions (OH⁻) in solution. ### Brønsted-Lowry Theory - A **Brønsted-Lowry acid** is a proton donor. - A **Brønsted-Lowry base** is a proton acceptor. ### Lewis Theory - A **Lewis acid** is an electron pair acceptor. - A **Lewis base** is an electron pair donor. ## pH Scale - A logarithmic scale used to express the acidity or basicity of a solution. - The pH scale ranges from 0 to 14, with 7 being neutral. - A pH less than 7 indicates an acidic solution, while a pH greater than 7 indicates a basic solution. ## Indicators - Substances that change color in the presence of an acid or a base. - They are used to determine the pH of a solution. ## Acid-Base Titration - A quantitative technique used to determine the concentration of an unknown acid or base solution using a solution of a known concentration as the titrant. ## Titration Curve - A graph that plots the pH of the solution being titrated against the volume of titrant added. ## Buffer Solutions - Solutions that resist changes in pH upon the addition of small amounts of acid or base. - They are composed of a weak acid and its conjugate base or a weak base and its conjugate acid. - They are essential in biological systems and many chemical processes. ## Solubility Product Constant (KsP) - A measure of the solubility of a sparingly soluble ionic compound. - It is the product of the concentrations of the ions in a saturated solution of the compound. ## Hydrolysis - The reaction of a salt with water to produce an acidic or basic solution. - Hydrolysis occurs because the ions of the salt can react with water molecules to form acidic or basic species. ## Oxidation-Reduction Reactions (Redox reactions) - Reactions that involve the transfer of electrons. - The species that loses electrons is oxidized while the species that gains electrons is reduced. ## Oxidation Numbers - A number assigned to an atom in a molecule or ion to indicate the number of electrons that atom has gained or lost. ## Balancing Redox Reactions - Redox reactions must be balanced to ensure that the number of atoms and charges on both sides of the equation are the same. ## Electrochemistry - The study of the relationship between chemical reactions and electrical energy. - It covers topics, such as: - Electrolysis - Electrochemical cells - Corrosion ## Electrolytic Cells - Electrochemical cells in which an external electric current is used to drive a non-spontaneous chemical reaction. ## Electrochemical Cells - Electrochemical cells that convert chemical energy into electrical energy. ## Corrosion - The deterioration of a metal surface due to chemical reactions with the environment. - Examples: rusting of iron, tarnishing of silver. ## Nuclear Chemistry - The study of the structure, properties, and reactions of atomic nuclei. - It covers topics such as: - Radioactivity - Nuclear fission - Nuclear fusion ## Radioactivity - The spontaneous emission of particles and/or energy from an unstable atomic nucleus is known as radioactivity. ## Nuclear Fission - The process of splitting a heavy atomic nucleus into lighter nuclei, releasing a tremendous amount of energy. ## Nuclear Fusion - The process of combining two or more lighter atomic nuclei to form a heavier nucleus, releasing a massive amount of energy.

Atomic Structure and Atomic Properties PDF

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