Electron Configuration and Chemical Bonding
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This document discusses electron configuration and chemical bonding. It begins by reviewing general ordering rules for electron configuration and defining core and valence electrons. It then discusses trends in atomic size and effective nuclear charge across the periodic table. As atomic number increases within a period, atomic radius generally decreases due to increased nuclear attraction. Transition metals are an exception, as their atomic radii remain relatively constant as valence electrons do not change.
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The document summarizes the historical development of the periodic table. It discusses early classification attempts by scientists like Dobereiner, Newlands, and Meyer. Mendeleev organized the 63 known elements into the first recognizable periodic table in 1869 based on periodic trends in their properties. Moseley's work on X-ray spectra led to the modern periodic table being arranged by atomic number instead of atomic mass. The periodic table is now understood to result from the periodic repetition of electronic structures as electrons fill different atomic orbitals in shells.
Classification of elements.pptx
The document discusses electronic configuration and periodic trends. It defines electronic configuration as the distribution of electrons in atomic orbitals. It then provides examples of the electronic configurations of hydrogen, oxygen, and chlorine. It also summarizes that properties like atomic radius, ionization energy, and electronegativity follow trends across periods and down groups as a result of electronic configuration.
Atomic+Structure (1).pptx
This document discusses the structure of atoms and bonding between atoms. It describes that an atom consists of a nucleus containing protons and neutrons, surrounded by electrons. Atoms bond through either ionic or covalent bonding. Ionic bonding occurs when atoms gain or lose electrons to become ions, forming electrostatic attractions. Covalent bonding occurs when atoms share electrons to achieve stable electron configurations.
EE2317-Course- 02 Atomic Structures and Interatomic bonding.pdf
This document discusses atomic structures and interatomic bonding. It begins by outlining the objectives and providing an outline of topics to be covered, including atomic models, electron configurations, the periodic table, and primary bonding types. It then delves into details of atomic structure including atomic number, mass, isotopes, and electron arrangements. The key bonding types of ionic, covalent and metallic are introduced and examples of each are provided. Ionic bonding involves transfer of electrons between metals and nonmetals, covalent bonding the sharing of electrons, and metallic bonding the sea of electrons in metal atoms.
Periodic trends
The document discusses periodic trends in elemental properties. It explains that Dmitri Mendeleev was the first to organize elements in a periodic table based on their properties. Elements in the same group have similar properties due to their valence electrons. Atomic radius generally decreases moving left to right across a period and increases moving down a group due to electron shielding. Ionization energy increases as atomic radius decreases. Electron affinity is exothermic when gaining electrons fills an orbital. Metallic character decreases and electronegativity increases moving from left to right. Cations are smaller than their parent atoms while anions are larger.
Periodic Trends
Dimitri Mendeleev was the first to publish an organized periodic table of elements in 1869. He understood the periodic law - that elements arranged by atomic number display a repeating pattern of chemical and physical properties in vertical columns due to their valence electrons. Mendeleev even predicted properties of undiscovered elements, which were later confirmed. The periodic table shows trends in atomic radius, ionization energy, electron affinity, and electronegativity that relate to chemical reactivity. Metals are on the left and are characterized by properties like conductivity, while nonmetals are on the right and have opposing properties.
electrochemistryclass12-180412053250.pptx
Electrochemistry is the study of chemical reactions involving the transfer of electrons. Oxidation and reduction reactions occur in electrochemical cells. Daniel cell is an example of a galvanic cell that converts chemical energy to electrical energy. It consists of zinc and copper half cells separated by a salt bridge. The cell potential depends on the standard electrode potentials of the half reactions and can be calculated using Nernst's equation. Equilibrium constants can also be determined from standard cell potentials using thermodynamic relationships.
ATOMS.pptx
Atoms are the smallest particles that can exist independently and are made up of subatomic particles like protons, neutrons, and electrons. Atoms of the same element have the same number of protons but can differ in the number of neutrons forming isotopes. Chemical bonds like ionic bonds form when electrons are transferred, covalent bonds when electrons are shared, and metallic bonds by the flow of electrons in metal atoms. The periodic table organizes elements according to their atomic structure.
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Ch. 7 electron configuration
- Electron configuration describes the arrangement of electrons in an atom's orbitals and energy levels according to 4 quantum numbers.
- The Pauli Exclusion Principle states that no two electrons can have the same set of 4 quantum numbers.
- Elements are arranged on the periodic table based on their electron configurations, with groups reflecting which orbitals are being filled.
- A noble gas configuration can be used to abbreviate long electron configurations by writing the nearest noble gas followed by the remaining electrons.
Periodic Relationships
There are observable trends in the physical and chemical properties of elements across periods and down groups in the periodic table. As you move across a period, atomic radius decreases due to the increasing effective nuclear charge pulling the valence electrons closer. Down groups, atomic radius increases as each successive energy level is farther from the nucleus. Ionization energy generally increases across periods as effective nuclear charge increases, and decreases down groups as the valence electrons are farther from the nucleus. Electron affinity increases across periods but decreases down groups. Oxidation states can be used to keep track of electrons in compounds and ions.
Lecture 1 medical Chemistry.pptxxfjxfjfjcf
This document provides an overview of the topics covered in the first lecture of a Medical Chemistry course. The lecture introduced the electronic structure of atoms, including the dual nature of electrons and quantum mechanical descriptions of hydrogen atoms. It also covered quantum numbers, atomic orbitals, electron configuration, and shielding effects in multi-electron atoms. The accompanying seminar discussed principal, angular momentum, magnetic, and electron spin quantum numbers as well as s, p, d, and f orbitals and their energies.
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Electrochemistry class 12 ( a continuation of redox reaction of grade 11)
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Electron Configuration and Chemical Bonding
- 1. Electron Configuration Chemical Bonding Amalene Cooper-Morgan, Ph.D.
- 2. General Ordering for Multi Electron Atoms
- 3. Recall… • Remember • S sublevel holds 2 electrons • P sublevel holds 6 electrons • D sublevel holds 10 electrons • F sublevel holds 14 electrons
- 4. Electron Configuration for Bromine
- 5. Use Noble Gases to Represent Inner Electrons
- 6. Valence Electrons • Located in the outermost principal energy level • For transition elements, d electrons are counted among the valence electrons, although they aren’t in the outermost principal energy level • Core electrons are electrons in complete principal energy levels
- 7. Valence vs Core Electrons
- 8. Main Group Elements • Don’t include transition elements • Valence electrons are in their outermost principal energy level • Example: • Germanium, atomic number 32 is a main group element • Ge n=1, 2 and 3 principal energy levels are complete (Core electrons) • Ge n=4 is the outermost principal level (Valence electrons)
- 11. Orbital blocks of the Periodic Table Except for the transition metals, main group elementss valence electrons are equal to the lettered group e.g Chlorine has 7 valence electrons and it’s in group 7A
- 12. Periodic Trends- Size of atoms and Effective Nuclear Charge • Orbitals do not represent a physical boundary but a statistical probability distribution of where to find an electron • Non bonding atomic radius or van der Waals radius- is the radius of an atom when it is not bonded to another atom
- 13. Periodic Trends- Size of atoms and Effective Nuclear Charge • Bonding atomic radius or covalent radius • Non-metals- ½ the distance between 2 of the atoms bonded together • Metals- ½ the distance between 2 of the atoms that are next to each other in a crystal of a metal
- 14. Exceptions.. • Bonding radii of He and Ne are approximated • They do not form either chemical bonds or metallic crystals
- 15. Atomic Radius • Average bonding radii • Radius of an atom when its bonded to another atom • Always smaller than van der waals radius
- 16. Approximate bond length • Sum of the atomic radii for 2 covalently bonded atoms • Approx. bond length for ICl: • I atomic radius = 133 pm + Cl atomic radius = 99 pm • Total bond length = 232 pm • Actual bond length is 232.07 pm
- 17. Trends across periodic table • Atomic radii peak with Alkali metals • Atomic radii is mainly determined by the valence electron • As we move • down a column, atomic radius decreases
- 18. Effective Nuclear Charge or Net Charge
- 19. Screening and Effective Nuclear Charge Core electrons = 1s2 electron Nucleus = 3+ Zeff ~ (3+) – 2 ~ 1+
- 20. Trends across the Periodic Table for Effective Nuclear Charge
- 21. Trends across periodic table • As we move across a row to the right the atomic radius decreases • Transition elements, except the first couple of elements in a column, do NOT follow the same trend for atomic radii as main group elements • Atomic radius does not decrease as you cross the row from L to R • Atomic radii stay constant • The outermost electrons stay constant and experience a constant nuclear charge, keeping the radius approx. constant