1.1 Atomic structure

Atomic structure forms the foundation of nuclear physics, providing insights into subatomic particles' behavior. From protons and neutrons in the nucleus to electrons orbiting in shells, understanding these components is crucial for manipulating matter at the atomic level.

Electron configuration, energy levels, and quantum principles like the Aufbau principle and Pauli exclusion principle explain atomic behavior. These concepts, along with atomic models, size, and electronic properties, help us grasp how atoms interact and form the basis for nuclear reactions and applications.

Fundamental atomic components

• Atomic structure forms the foundation of nuclear physics, providing insights into the behavior of subatomic particles
• Understanding fundamental atomic components enables the manipulation of matter at the atomic level, crucial for nuclear applications

Protons, neutrons, electrons

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• Protons carry a positive charge and reside in the nucleus
• Neutrons lack electrical charge and coexist with protons in the nucleus
• Electrons possess a negative charge and orbit the nucleus in electron shells
• Proton and neutron masses approximately equal, while electron mass ~1/1836 of a proton
• Strong nuclear force binds protons and neutrons within the nucleus

Atomic number vs mass number

• Atomic number (Z) represents the number of protons in an atom's nucleus
• Mass number (A) indicates the total number of protons and neutrons
• Relationship between atomic and mass numbers: A = Z + N (where N is the number of neutrons)
• Determines an element's position in the periodic table and its chemical properties
• Mass number influences isotopic variations and nuclear stability

Isotopes and nuclides

• Isotopes are atoms of the same element with different numbers of neutrons
• Nuclides refer to specific combinations of protons and neutrons in atomic nuclei
• Notation for nuclides: $^A_Z X$ (X represents the element symbol)
• Isotopes exhibit identical chemical properties but different physical properties
• Some isotopes are stable, while others undergo radioactive decay (radioisotopes)

Electron configuration

• Electron configuration describes the arrangement of electrons in an atom's orbitals
• Crucial for understanding chemical bonding, spectroscopy, and atomic behavior in nuclear reactions

Energy levels and orbitals

• Electrons occupy discrete energy levels or shells around the nucleus
• Principal quantum number (n) defines the main energy levels (1, 2, 3, etc.)
• Sublevels or orbitals within each energy level (s, p, d, f)
• Orbitals represent regions of high electron probability
• Energy increases as electrons occupy higher energy levels

Aufbau principle

• Electrons fill orbitals from lowest to highest energy levels
• Order of orbital filling: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p
• Determines the ground state electron configuration of atoms
• Explains periodic trends in atomic properties
• Exceptions occur due to stability of half-filled or fully-filled subshells

Hund's rule

• Electrons occupy orbitals of equal energy individually before pairing
• Maximizes total spin angular momentum of the atom
• Results in the most stable electron configuration
• Applies to p, d, and f orbitals with multiple sublevels
• Influences magnetic properties of atoms and molecules

Pauli exclusion principle

• No two electrons in an atom can have identical quantum numbers
• Limits the number of electrons in each orbital to two
• Electrons in the same orbital must have opposite spins
• Fundamental principle in quantum mechanics
• Explains electron shell structure and periodic table organization

Atomic models

• Atomic models evolved to explain observed atomic behavior and properties
• Each model contributed to our current understanding of atomic structure in nuclear physics

Dalton's model

• Proposed indivisible, indestructible atoms as basic units of matter
• Explained chemical reactions through combinations of atoms
• Introduced the concept of atomic weight
• Limited by lack of knowledge about subatomic particles
• Laid foundation for modern atomic theory

Thomson's plum pudding model

• Depicted atoms as positively charged "pudding" with embedded negative electrons
• Explained the neutrality of atoms
• Proposed electrons could be removed from atoms (ionization)
• Failed to account for the nucleus or atomic mass distribution
• Led to the discovery of the electron

Rutherford's nuclear model

• Introduced the concept of a dense, positively charged nucleus
• Based on the famous gold foil experiment
• Explained the deflection of alpha particles by thin metal foils
• Proposed electrons orbit the nucleus like planets around the sun
• Unable to explain atomic spectra or electron stability

Bohr's model

• Introduced quantized electron energy levels
• Explained hydrogen atom spectrum and atomic stability
• Proposed electrons move in circular orbits around the nucleus
• Introduced the concept of ground state and excited states
• Limited to hydrogen-like atoms and failed for multi-electron systems

Quantum mechanical model

• Describes electrons as wave functions in probability clouds
• Introduces quantum numbers to characterize electron states
• Explains multi-electron atoms and molecular bonding
• Accounts for the uncertainty principle and wave-particle duality
• Forms the basis of modern understanding of atomic structure

Atomic size

• Atomic size influences chemical reactivity, bonding, and physical properties
• Critical for understanding atomic interactions in nuclear processes

Atomic radius

• Measures the distance from the nucleus to the outermost electron shell
• Decreases across a period due to increasing nuclear charge
• Increases down a group due to additional electron shells
• Affected by electron-electron repulsion and nuclear attraction
• Influences atomic volume and density

Ionic radius

• Size of an atom after gaining or losing electrons to form an ion
• Cations (positively charged ions) smaller than parent atoms
• Anions (negatively charged ions) larger than parent atoms
• Depends on the charge of the ion and its electron configuration
• Affects ionic bonding and crystal structure formation

Periodic trends

• Atomic radius generally decreases across a period and increases down a group
• Ionization energy increases across a period and decreases down a group
• Electron affinity generally increases across a period and decreases down a group
• Electronegativity increases across a period and decreases down a group
• Metallic character decreases across a period and increases down a group

Electronic properties

• Electronic properties determine an atom's behavior in chemical reactions and nuclear processes
• Essential for understanding atomic interactions and bonding mechanisms

Ionization energy

• Energy required to remove an electron from a neutral atom in its ground state
• First ionization energy removes the outermost electron
• Subsequent ionization energies progressively higher due to increased nuclear attraction
• Increases across a period and decreases down a group
• Relates to an element's reactivity and ability to form cations

Electron affinity

• Energy change when a neutral atom gains an electron
• Exothermic process for most elements (energy released)
• Generally increases across a period and decreases down a group
• Highest for halogens, lowest for noble gases
• Influences an element's tendency to form anions

Electronegativity

• Measure of an atom's ability to attract shared electrons in a chemical bond
• Increases across a period and decreases down a group
• Fluorine has the highest electronegativity, francium the lowest
• Determines bond polarity and molecular dipole moments
• Pauling scale most commonly used for electronegativity values

Atomic spectra

• Atomic spectra provide insights into electron energy levels and transitions
• Crucial for spectroscopic analysis and understanding atomic structure

Emission vs absorption spectra

• Emission spectra produced when excited atoms release energy as photons
• Absorption spectra result from atoms absorbing specific wavelengths of light
• Emission spectra appear as bright lines on a dark background
• Absorption spectra show dark lines on a continuous spectrum
• Both types unique to each element, serving as "fingerprints"

Spectral lines

• Discrete lines in atomic spectra correspond to specific electron transitions
• Wavelengths of spectral lines relate to energy differences between electron states
• Balmer series in hydrogen spectrum represents transitions to n=2 energy level
• Fine structure results from spin-orbit coupling in multi-electron atoms
• Hyperfine structure arises from interactions between electrons and nuclear spin

Quantum transitions

• Electrons move between energy levels through absorption or emission of photons
• Energy of emitted or absorbed photon: $E = hf = (E_2 - E_1)$
• Selection rules govern allowed transitions between quantum states
• Spontaneous emission occurs without external stimulation
• Stimulated emission forms the basis for laser operation

Nuclear structure

• Nuclear structure determines isotope stability and radioactive decay processes
• Fundamental to understanding nuclear reactions and energy production

Nuclear force

• Strong nuclear force binds protons and neutrons in the nucleus
• Overcomes electrostatic repulsion between protons at short distances
• Short-range force, effective only within ~1 femtometer
• Charge-independent, acts equally on protons and neutrons
• Responsible for nuclear stability and binding energy

Binding energy

• Energy required to separate a nucleus into its constituent nucleons
• Calculated using mass defect: $\Delta E = \Delta m c^2$
• Binding energy per nucleon peaks around iron (most stable nuclei)
• Explains energy release in nuclear fission and fusion reactions
• Determines nuclear stability and susceptibility to radioactive decay

Nuclear stability

• Stable nuclei have optimal neutron-to-proton ratios
• Light elements stable with equal numbers of protons and neutrons
• Heavier elements require more neutrons for stability (neutron excess)
• Nuclear shell model explains "magic numbers" of exceptionally stable nuclei
• Unstable nuclei undergo radioactive decay to achieve more stable configurations

Atomic mass

• Atomic mass concepts crucial for understanding isotopic compositions and nuclear reactions
• Fundamental to calculations in nuclear physics and chemistry

Atomic mass unit

• Defined as 1/12 the mass of a carbon-12 atom
• Approximately equal to the mass of a proton or neutron
• Used to express masses of atoms and subatomic particles
• 1 amu = 1.660539 × 10^-27 kg
• Convenient unit for atomic-scale mass measurements

Relative atomic mass

• Average mass of an element's atoms relative to 1/12 the mass of carbon-12
• Accounts for naturally occurring isotopes and their abundances
• Calculated as weighted average of isotopic masses
• Appears on the periodic table as the atomic weight
• Used in stoichiometric calculations and determining molecular masses

Isotopic abundance

• Relative amounts of different isotopes of an element in nature
• Expressed as percentages or fractions of total atoms
• Varies by geographical location and can change over time
• Determined using mass spectrometry techniques
• Influences the average atomic mass of elements

Quantum numbers

• Quantum numbers describe the state of electrons in atoms
• Essential for understanding electron configuration and spectroscopic transitions

Principal quantum number

• Denoted by n, represents the main energy level of an electron
• Takes positive integer values (1, 2, 3, ...)
• Determines the overall energy and size of the orbital
• Relates to the radial probability distribution of electrons
• Higher n values correspond to higher energy and larger orbitals

Azimuthal quantum number

• Represented by l, describes the shape of the orbital
• Takes values from 0 to (n-1)
• Corresponds to subshells s (l=0), p (l=1), d (l=2), f (l=3)
• Determines the orbital angular momentum of the electron
• Influences the magnetic properties of atoms

Magnetic quantum number

• Denoted by m_l, specifies the orientation of the orbital in space
• Takes integer values from -l to +l, including zero
• Determines the number of orbitals within a subshell
• Relates to the z-component of orbital angular momentum
• Important in the Zeeman effect and magnetic interactions

Spin quantum number

• Represented by m_s, describes the intrinsic angular momentum of an electron
• Takes values of +1/2 or -1/2
• Arises from the electron's inherent magnetic moment
• Explains fine structure in atomic spectra
• Crucial for understanding the Pauli exclusion principle and electron pairing

Atomic interactions

• Atomic interactions govern chemical bonding and physical properties of matter
• Understanding these interactions is crucial for predicting nuclear behavior

Chemical bonding basics

• Covalent bonds involve sharing of electrons between atoms
• Ionic bonds result from electrostatic attraction between oppositely charged ions
• Metallic bonds occur in metals, with delocalized electrons in a "sea"
• Bond strength influences physical properties (melting point, boiling point)
• Molecular geometry determined by electron pair repulsion theory

Intermolecular forces

• Van der Waals forces include dipole-dipole, dipole-induced dipole, and London dispersion forces
• Hydrogen bonding occurs between H atoms and electronegative atoms (N, O, F)
• Ion-dipole interactions important in solutions of ionic compounds
• Pi-stacking influences structure of aromatic compounds
• Determine physical properties like boiling point and solubility

Atomic excitation and relaxation

• Excitation occurs when atoms absorb energy, promoting electrons to higher energy levels
• Relaxation involves electrons returning to lower energy states, emitting photons
• Fluorescence results from rapid relaxation with visible light emission
• Phosphorescence involves longer-lived excited states and delayed emission
• Crucial for understanding atomic spectra and laser operation