ATOMS (1).pdf

ATOMS 2 AGENDA Introduction Atomic Theory Sub-atomic Particles, mass number and atomic number Quantum Numbers Electron Distribution Periodic Table ATOMS ATOMS Chemistry has been practiced for a very long time, even if its practitioners were much more interested in its applications than in its underlying principles. The blast furnace for extracting iron from iron ore appeared as early as A.D. 1300, and such important chemicals as sulfuric acid (oil of vitriol), nitric acid (aqua fortis), and sodium sulfate (Glauber s salt) were all well known and used several hundred years ago. Before the end of the eighteenth century, the principal gases of the atmosphere nitrogen and oxygen had been isolated, and natural laws had been proposed describing the physical behavior of gases. Yet chemistry cannot be said to have entered the modern age until the process of combustion was explained. ATOMS Democritus: Matter composed of tiny particles called “Atomos” Atomos is a Greek word, meaning (not to be cut or to be divided. 6 ATOMOS Átomos" (ἄτομος) is a term from ancient Greek philosophy meaning "indivisible" or "uncuttable." The concept of the atom was first introduced by the Greek philosopher Democritus around 400 BCE. Democritus and his teacher Leucippus proposed that all matter is composed of small, indivisible particles called "atoms." Key ideas from Democritus' atomic theory: 1. Atoms are indivisible: Democritus believed that atoms were the smallest unit of matter, and they could not be divided further. They were eternal and indestructible. 2. Atoms differ in shape and size: According to Democritus, different types of atoms had different shapes and sizes, which explained the diversity of materials in the natural world. 3. Atoms are in constant motion: He proposed that atoms move in a void (empty space) and collide with each other, leading to the formation of different substances. 4. Matter and void: Democritus' theory suggested that the universe is made up of two fundamental things: atoms and the void (empty space in which atoms move). 7 ATOMS Atoms are the fundamental building blocks of matter, consisting of a central nucleus surrounded by electrons. Structure of an Atom 1. Nucleus: 1. Protons: Positively charged particles found in the nucleus. The number of protons in an atom’s nucleus determines the element's atomic number and identity. 2. Neutrons: Neutral particles (with no charge) found in the nucleus. Neutrons, along with protons, contribute to the atomic mass but do not affect the element's chemical properties. 2. Electrons: 1. Electrons: Negatively charged particles that orbit the nucleus in various energy levels or electron shells. The arrangement of electrons determines the atom's chemical behavior and bonding properties. 8 ATOMIC STRUCTURE Nucleus: The nucleus is the dense, central core of the atom. It contains two types of subatomic particles: Protons: Positively charged particles. The number of protons in an atom determines the atomic number (Z) and defines the element. For example, hydrogen has 1 proton, oxygen has 8, and gold has 79. Neutrons: Neutral particles (no charge). The number of neutrons can vary between atoms of the same element, resulting in different isotopes of that element. The nucleus is tiny compared to the overall size of the atom but contains nearly all of its mass. 9 ATOMIC STRUCTURE Electrons: Electrons are negatively charged particles that move around the nucleus in regions of space called orbitals. Unlike protons and neutrons, electrons have very little mass (about 1/1836th the mass of a proton), but they are crucial to chemical bonding and reactions. Electrons are organized into energy levels or shells, and within these levels, they occupy specific orbitals (s, p, d, f). 10 SUBATOMIC PARTICLES Subatomic Particle Charge Mass Location 1 atomic mass unit Proton +1 Nucleus (amu) Neutron 0 (neutral) 1 amu Nucleus Outside the nucleus in Electron -1 ~0.0005 amu orbitals 11 IMPORTANT CONCEPTS IN ATOMIC STRUCTURE 1. Atomic Number (Z): The atomic number is the number of protons in the nucleus of an atom. It determines the identity of the element and its position on the periodic table. Example: The atomic number of carbon is 6, meaning every carbon atom has 6 protons. 12 2. MASS NUMBER The mass number is the total number of protons and neutrons in the nucleus of an atom. It is approximately equal to the atomic mass of the atom. Example: For an isotope of carbon, carbon-12 has 6 protons and 6 neutrons, giving it a mass number of 12. 13 3. ISOTOPES Isotopes are atoms of the same element (same number of protons) but with different numbers of neutrons. Example: Carbon-12 (⁶C¹²) and Carbon-14 (⁶C¹⁴) are isotopes of carbon, with 6 and 8 neutrons, respectively. They have different mass numbers but the same atomic number. 14 ATOMIC SPECIES Similarities Difference Isotopes Element Mass Number Isotones Neutrons Protons Isobars Mass number Element Isomers Molecule Structure 15 4. ELECTRON CONFIGURATION Electrons are arranged in energy levels (shells) around the nucleus, and within these shells, they occupy orbitals of specific shapes and energies. The arrangement of electrons, or electron configuration, determines an element’s chemical properties, including its reactivity and bonding behavior. 16 5. ENERGY LEVELS AND ORBITALS Electrons occupy specific energy levels denoted by the principal quantum number n (n = 1, 2, 3, etc.). These energy levels are subdivided into sublevels: s, p, d, and f, each with a specific shape and number of orbitals. s orbitals: Spherical, 1 orbital per sublevel. p orbitals: Dumbbell-shaped, 3 orbitals per sublevel. d orbitals: Complex shapes, 5 orbitals per sublevel. f orbitals: Even more complex, 7 orbitals per sublevel. The maximum number of electrons each energy level can hold is determined by the formula 2n². 17 18 6. VALENCE ELECTRON Valence electrons are the electrons in the outermost energy level of an atom. They play a crucial role in chemical bonding and determining how an atom interacts with others. Atoms with similar numbers of valence electrons are grouped together in the periodic table because they have similar chemical properties. 19 7. ATOMIC MASS The atomic mass (also known as atomic weight) is the weighted average mass of an element’s isotopes as they occur naturally. It takes into account the relative abundance of each isotope. Atomic mass is typically close to the mass number, but it can vary slightly due to the presence of different isotopes. 20 ATOMIC THEORY 21 DALTONS ATOMIC THEORY Dalton's Atomic Theory was one of the first modern scientific theories of atoms, proposed by the English chemist John Dalton in 1808. It laid the foundation for our understanding of matter and chemical reactions. The key postulates of Dalton's atomic theory are: 22 DALTONS ATOMIC THEORY All matter is made up of tiny, indivisible particles called atoms: Dalton proposed that atoms are the smallest unit of matter and cannot be divided or destroyed. This was later revised as we learned about subatomic particles like electrons, protons, and neutrons. Atoms of the same element are identical in mass and properties: Atoms of a given element have the same size, mass, and chemical properties. However, this idea was later modified when isotopes (atoms of the same element with different masses) were discovered. Atoms of different elements differ in mass and properties: Dalton suggested that atoms of different elements are unique and differ from one another in their mass and other characteristics. Atoms combine in fixed ratios to form compounds: Compounds are composed of atoms from different elements combined in simple, whole- number ratios. This explains why compounds always have the same proportions of elements. In chemical reactions, atoms are rearranged, but not created or destroyed: Dalton's theory adhered to the law of conservation of mass, which states that atoms in a chemical reaction are only rearranged to form new compounds, but the total number of atoms remains the same. 23 DISCOVERY OF SUBATOMIC PARTICLES J.J. Thomson (1897) - Joseph John Thompson discovered the electron using cathode ray experiments. He proposed the plum pudding model, where negatively charged electrons are embedded in a positively charged sphere. 24 25 KEY FEATURES OF THE PLUM PUDDING MODEL Positively charged "pudding": Thomson suggested that atoms were composed of a large, positively charged "soup" or "pudding" that spread throughout the entire atom, which accounted for the atom’s overall positive charge. Electrons embedded like "plums": Negatively charged electrons were thought to be scattered or embedded within this positive "pudding" or sphere, like "plums" in a traditional British plum pudding (a dessert). The electrons were small compared to the positively charged matter. Neutral overall charge: Since the atom is electrically neutral, the positive charge of the "pudding" exactly balanced the negative charge of the electrons. 26 NUCLEAR MODEL Ernest Rutherford (1911) discovered the nucleus through his gold foil experiment, which showed that an atom has a small, dense, positively charged nucleus, and the rest is mostly empty space. He proposed the nuclear model of the atom. 27 GOLD FOIL EXPERIMENT Rutherford's conclusion came from his experiment where he directed a beam of alpha particles (positively charged) at a very thin sheet of gold foil. According to the Plum Pudding Model, these particles should have passed through with only minor deflections. Instead, Rutherford observed that while most particles passed through, a small fraction were deflected at large angles, and some even bounced back. This surprising result led to the conclusion that the positive charge in the atom was not spread throughout, as in the Plum Pudding Model, but concentrated in a tiny central region—the nucleus. 28 KEY FEATURES OF RUTHERFORD’S NUCLEAR MODEL Nucleus at the center: The atom consists of a tiny, dense, positively charged core called the nucleus. The nucleus contains most of the atom's mass, and it is made up of protons (positively charged particles) and, as later discovered, neutrons (neutral particles). Electrons orbit the nucleus: Negatively charged electrons orbit the nucleus at a relatively large distance compared to the size of the nucleus. The space between the nucleus and the electrons is mostly empty, explaining why most alpha particles in the gold foil experiment passed through the foil undisturbed. Atom is mostly empty space: Most of the atom's volume is empty space. This explained why only a small fraction of alpha particles were deflected in the gold foil experiment, while most passed straight through. Nuclear size vs. atomic size: The nucleus is extremely small compared to the overall size of the atom. Rutherford’s model revealed that the atom's positive charge and nearly all its mass are concentrated in this very small central region. 29 BOHR MODEL Niels Bohr (1913) improved upon Rutherford’s model by introducing the idea that electrons orbit the nucleus in specific energy levels, or shells, leading to the Bohr model of the atom. 30 KEY FEATURES OF BOHR MODEL Electrons orbit the nucleus in fixed energy levels (shells): Bohr proposed that electrons move around the nucleus in specific, fixed orbits or energy levels, much like planets orbiting the sun. These orbits are at certain allowed distances from the nucleus, and each corresponds to a specific energy level. Quantized energy levels: Unlike in classical physics, where an electron could theoretically lose energy and spiral into the nucleus, Bohr suggested that electrons can only occupy certain discrete energy levels. The energy of an electron is "quantized," meaning it can only exist in specific amounts corresponding to its orbit. Electrons can transition between energy levels: Electrons can move from one energy level to another, but they cannot exist between levels. When an electron jumps from a higher energy level (excited state) to a lower one (ground state), it releases energy in the form of light (a photon). The energy of the photon corresponds to the difference in energy between the two levels. This concept explained the line spectra of elements, particularly hydrogen. Energy absorption and emission: When an electron absorbs energy (e.g., from heat or light), it can move to a higher energy level (excited state). When it falls back to a lower energy level, it emits energy in the form of light (this emission of light creates the characteristic line spectra of elements). 31 QUANTUM MECHANICAL MODEL OF THE ATOM (1920S AND BEYOND) 32 QUANTUM MECHANICAL MODEL OF THE ATOM The development of quantum mechanics led to a more accurate understanding of atomic structure. Erwin Schrödinger introduced the idea that electrons are not in fixed orbits but exist in probability clouds called orbitals. This is known as the quantum mechanical model of the atom. Werner Heisenberg's Uncertainty Principle stated that it is impossible to precisely know both the position and momentum of an electron at the same time. James Chadwick (1932) discovered the neutron, completing the understanding of the atomic nucleus as being composed of protons and neutrons. KEY FEATURES OF QUANTUM 33 MECHANICAL MODEL OF THE ATOM Electrons exist in orbitals, not fixed paths: Unlike Bohr’s model, where electrons travel in fixed orbits around the nucleus, the quantum mechanical model describes electrons as existing in regions of space called orbitals. These are probability distributions that represent where an electron is likely to be found, rather than precise orbits. An orbital is a three-dimensional region around the nucleus where there is a high probability (usually 90%) of finding an electron. Wave-particle duality of electrons: Electrons exhibit both particle-like and wave-like properties. Louis de Broglie suggested that electrons could behave as waves, and Schrödinger built on this with his famous wave equation, which describes the behavior of electrons as standing waves in the atom. This wave-particle duality is a fundamental concept in quantum mechanics. Heisenberg’s Uncertainty Principle: Werner Heisenberg showed that it is impossible to precisely determine both the position and momentum of an electron at the same time. This is known as the Uncertainty Principle. Because of this, we can only talk about the probability of finding an electron in a certain region of space, rather than a fixed path. KEY FEATURES OF QUANTUM 34 MECHANICAL MODEL OF THE ATOM Schrödinger’s Wave Equation: Erwin Schrödinger formulated the Schrödinger equation, a mathematical equation that describes the wave-like behavior of electrons. Solutions to this equation, called wavefunctions (denoted by the Greek letter ψ), provide information about the probability distribution of an electron in an atom. The square of the wavefunction (ψ²) gives the probability density, telling us where the electron is most likely to be found. Quantum numbers: The behavior of electrons is described by quantum numbers, which arise from the solutions to Schrödinger's equation. These quantum numbers describe various properties of electrons, such as their energy level, shape of the orbital, orientation, and spin. The four quantum numbers are: 1.Principal quantum number (n): Determines the energy level of the electron and the size of the orbital. 2.Angular momentum quantum number (l): Defines the shape of the orbital (s, p, d, f orbitals). 3.Magnetic quantum number (mₗ): Describes the orientation of the orbital in space. 4.Spin quantum number (mₛ): Describes the spin of the electron, which can be +½ or - ½. 35 QUANTUM NUMBERS l (subshell) Example: s= 0 1. F 2p5 p=1 - n=2 d=2 f=3 - l=1 ml correspond to certain order - ml= 0 (hunds rule) s=0 p= -1, 0, 1 d= -2, -1, 0, 1, 2 f= -3, -2, -1, 0, 1, 2, 3 ms Arrow down -1/2 Arrow up +1/2 KEY FEATURES OF QUANTUM 36 MECHANICAL MODEL OF THE ATOM Orbitals and their shapes: The shape and size of an electron’s orbital depend on its energy and the quantum numbers. The most common orbitals are: s-orbitals: Spherical in shape. p-orbitals: Dumbbell-shaped. d-orbitals and f-orbitals: More complex shapes. These orbitals represent regions where electrons are most likely to be found. ELECTRON DISTRIBUTION 38 ELECTRON CONFIGURATION Electron configuration is a notation that shows the distribution of electrons among the orbitals of an atom. Electrons fill orbitals in order of increasing energy according to the Aufbau principle, Hund's rule, and the Pauli Exclusion Principle: Aufbau principle: Electrons occupy the lowest energy orbital available. Hund’s rule: Electrons will fill orbitals of the same sublevel singly before pairing. Pauli Exclusion Principle: No two electrons in the same atom can have the same set of four quantum numbers, so each orbital can hold a maximum of two electrons with opposite spins. 39 AUFBAU PRINCIPLE The Aufbau Principle (also called the building-up principle or the Aufbau rule) states that, in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy level before occupying higher-energy levels. 40 ACTIVITY: Identify the electron configuration of the given element: 1. H 2. He 3. O 4. Ca 5. Iron 41 ANSWER: Hydrogen (H): 1 electron → 1s¹ Helium (He): 2 electrons → 1s² Oxygen (O): 8 electrons → 1s² 2s² 2p⁴ Calcium (Ca): 20 electrons → 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² Iron (Fe): 26 electrons → 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s² 42 HUND’S RULE OF MAXIMUM MULTIPLICITY Hund’s rule: Electrons will fill orbitals of the same sublevel singly before pairing. 43 PAULI’S EXCLUSION Pauli Exclusion Principle: No two electrons in the same atom can have the same set of four quantum numbers, so each orbital can hold a maximum of two electrons with opposite spins. THE PERIODIC TABLE OF ELEMENTS 45 THE PERIODIC TABLE The periodic table of elements is a tabular arrangement of chemical elements, organized based on their atomic number, electron configuration, and recurring chemical properties. The structure of the periodic table allows scientists to predict the characteristics and behaviors of elements, making it a fundamental tool in chemistry. IMPORTANT INDIVIDUAL IN 47 THE DISCOVERY OF PERIODIC TABLE First List of Chemical Elements Antoine-Laurent Lavoisier (1789): (1743–1794) is often referred Lavoisier was one of the first to as the "Father of Modern scientists to classify substances into elements, those that could not be Chemistry" due to his broken down further by chemical revolutionary work in means. In his seminal work, "Traité understanding the nature of Élémentaire de Chimie" (Elementary Treatise of Chemistry), he identified chemical elements and the 33 substances he considered laws of chemical reactions. elements, including oxygen, While he did not directly nitrogen, hydrogen, phosphorus, mercury, zinc, and sulfur. contribute to the creation of the modern periodic table, Although not all of these substances were elements (some his work laid the foundation for were later found to be compounds), future discoveries and Lavoisier’s systematic classification classifications of elements. represented the first serious attempt to list and define elements scientifically. IMPORTANT INDIVIDUAL IN 48 THE DISCOVERY OF PERIODIC TABLE Dmitri Mendeleev (1834–1907): Julius Lothar Meyer (1830–1895): Contribution: Mendeleev is considered Contribution: Meyer independently the father of the modern periodic table. In developed a periodic table around the 1869, he arranged the known elements same time as Mendeleev. He also into a table based on their atomic masses organized elements based on atomic and observed recurring patterns, or mass and noticed the periodicity of periodic trends, in their properties. their properties. Key Achievement: Mendeleev's table was unique because he left gaps for Key Achievement: In 1864, Meyer undiscovered elements and correctly published an early version of the predicted their properties. For example, periodic table containing only 28 he predicted the existence and properties elements. In 1870, after Mendeleev’s of elements like gallium, scandium, and table, Meyer published a more germanium before they were discovered. comprehensive table, showing how atomic volume varied with atomic Legacy: Although some of his mass. assumptions (like ordering elements strictly by atomic mass) were later Legacy: Meyer is often considered co- modified, Mendeleev’s work laid the creator of the periodic table, but foundation for the modern periodic table. Mendeleev’s predictions of new elements gave him more recognition in the scientific community. IMPORTANT INDIVIDUAL IN 49 THE DISCOVERY OF PERIODIC TABLE Henry Moseley (1887–1915): John Newlands (1837–1898): Contribution: Moseley made a crucial Contribution: In 1864, Newlands modification to Mendeleev’s periodic table proposed the Law of Octaves, by organizing elements according to their atomic number (the number of protons in suggesting that elements could be an atom) rather than their atomic mass. arranged in "octaves" because every eighth element had similar Key Achievement: Using X-ray properties when arranged by spectroscopy, Moseley discovered that increasing atomic mass. each element has a unique atomic number, which resolved inconsistencies in Key Achievement: Although his Mendeleev's table. For example, iodine (I) law was not entirely accurate, and tellurium (Te) are correctly placed in order of atomic number rather than atomic Newlands was the first to recognize mass. a periodic pattern in the properties of elements, foreshadowing later Legacy: Moseley’s work established that developments. the atomic number, not atomic mass, is the correct way to organize the periodic Legacy: His work was initially table. His discovery led to the modern dismissed but later recognized as an arrangement of elements. important step toward understanding periodicity in elements. IMPORTANT INDIVIDUAL IN 50 THE DISCOVERY OF PERIODIC TABLE Johann Wolfgang Döbereiner (1780–1849): Glenn T. Seaborg (1912–1999): Developed the concept of triads, Contribution: Seaborg made significant contributions to the grouping elements with similar discovery of transuranium properties in sets of three. For elements (elements beyond uranium in the periodic table, atomic example, lithium, sodium, and numbers greater than 92). potassium formed a triad. His work Key Achievement: He laid the groundwork for later reconfigured the periodic table to developments in periodic include the actinide series, placing them below the lanthanides. This classification. modification is now a standard William Ramsay (1852–1916): feature of the modern periodic table. Discovered the noble gases (e.g., Legacy: Seaborg co-discovered 10 elements and won the Nobel Prize argon, neon, krypton, and xenon), in Chemistry in 1951. Element 106, which led to the addition of a new seaborgium (Sg), is named in his honor. group (Group 18) to the periodic table. 51 STRUCTURE OF PERIODIC TABLE Groups (Columns): Periods (Rows): The vertical columns in the periodic table are called The horizontal rows in the groups or families. periodic table are called There are 18 groups, numbered from 1 to 18. periods. Elements within the same group have similar chemical There are 7 periods, properties because they have the same number of corresponding to the number valence electrons (electrons in their outermost energy level). of electron shells or energy levels in the atoms of the Important groups include: elements. Group 1: Alkali Metals (e.g., Lithium, Sodium, Potassium) As you move from left to right across a period, the Group 2: Alkaline Earth Metals (e.g., atomic number increases, Magnesium, Calcium) and elements become less Group 17: Halogens (e.g., Fluorine, Chlorine) metallic in character. Group 18: Noble Gases (e.g., Helium, Neon, Argon) 52 STRUCTURE OF PERIODIC ELEMENT Metals, Nonmetals, and Element Blocks: Metalloids: The periodic table is divided into four blocks Metals: Found on the left and based on the subshell that is being filled with center of the periodic table (e.g., iron, gold, sodium). They are good electrons: conductors of heat and electricity, malleable, and ductile. s-block: Groups 1 and 2 (and hydrogen Nonmetals: Located on the right and helium). side of the periodic table (e.g., p-block: Groups 13 to 18. oxygen, carbon, nitrogen). They are poor conductors of heat and d-block: Transition metals (Groups 3 to electricity and are often brittle in solid form. 12). f-block: Lanthanides and actinides (rare Metalloids: Elements that have properties of both metals and earth elements, usually displayed nonmetals (e.g., silicon, arsenic). They are located along a diagonal separately at the bottom of the table). line between metals and nonmetals. 53 PERIODIC TRENDS Atomic Radius: Ionization Energy: Atomic size increases as you move down a group Ionization energy is the because additional electron shells are added. energy required to remove an electron from an atom. Atomic size decreases as you move across a period (left to right) because the increasing number It decreases down a group of protons pulls the electron cloud closer to the because electrons are nucleus. farther from the nucleus. It increases across a period because atoms have stronger attractions to their electrons due to the increasing number of protons. 54 PERIODIC TRENDS Electronegativity: Metallic Character: Electronegativity is the tendency of an atom to Metallic character refers to attract electrons in a chemical bond. how easily an element loses electrons to form positive It decreases down a group and increases across a ions. period. It increases down a group Fluorine is the most electronegative element, while and decreases across a francium is the least. period. THANK YOU

Related

Transcript

Upgrade to continue