﻿ Concept Summary - Atomic Structure

# Atomic StructureConcept Summary

Subatomic Particles

· A proton has a positive charge and mass around 1 amu; a neutron has no charge and mass around 1 amu; an electron has a negative charge and negligible mass.

· The nucleus contains the protons and neutrons, while the electrons move around the nucleus.

· The atomic number is the number of protons in a given element.

· The mass number is the sum of an element’s protons and neutrons.

Atomic Mass vs. Atomic Weight

· Atomic mass is essentially equal to the mass number, the sum of an element’s protons and neutrons.

o Isotopes are atoms of a given element (same atomic number) that have different mass numbers. They differ in the number of neutrons.

o Most isotopes are identified by the element followed by the mass number (such as carbon-12, carbon-13, and carbon-14).

o The three isotopes of hydrogen go by different names: protium, deuterium, and tritium.

· Atomic weight is the weighted average of the naturally occurring isotopes of an element. The periodic table lists atomic weights, not atomic masses.

Rutherford, Planck, and Bohr

· Rutherford first postulated that the atom had a dense, positively charged nucleus that made up only a small fraction of the volume of the atom.

· In the Bohr model of the atom, a dense, positively charged nucleus is surrounded by electrons revolving around the nucleus in orbits with distinct energy levels.

· The energy difference between energy levels is called a quantum, first described by Planck.

o Quantization means that there is not an infinite range of energy levels available to an electron; electrons can exist only at certain energy levels. The energy of an electron increases the farther it is from the nucleus.

o The atomic absorption spectrum of an element is unique; for an electron to jump from a lower energy level to a higher one, it must absorb an amount of energy precisely equal to the energy difference between the two levels.

o When electrons return from the excited state to the ground state, they emit an amount of energy that is exactly equal to the energy difference between the two levels; every element has a characteristic atomic emission spectrum, and sometimes the electromagnetic energy emitted corresponds to a frequency in the visible light range.

Quantum Mechanical Model of Atoms

· The quantum mechanical model posits that electrons do not travel in defined orbits but rather are localized in orbitals; an orbital is a region of space around the nucleus defined by the probability of finding an electron in that region of space.

· The Heisenberg uncertainty principle states that it is impossible to know both an electron’s position and its momentum exactly at the same time.

· There are four quantum numbers; these numbers completely describe any electron in an atom.

o The principal quantum number, n, describes the average energy of a shell.

o The azimuthal quantum number, l, describes the subshells within a given principal energy level (s, p, d, and f).

o The magnetic quantum number, ml, specifies the particular orbital within a subshell where an electron is likely to be found at a given moment in time.

o The spin quantum number, ms, indicates the spin orientation of an electron in an orbital.

· The electron configuration uses spectroscopic notation (combining the n and l values as a number and letter, respectively) to designate the location of electrons.

o For example, 1s22s22p63s2 is the electron configuration for magnesium: a neutral magnesium atom has 12 electrons—two in the s subshell of the first energy level, two in the s subshell of the second energy level, six in the p subshell of the second energy level, and two in the s subshell of the third energy level; the two electrons in the 3s subshell are the valence electrons for the magnesium atom.

· Electrons fill the principal energy levels and subshells according to increasing energy, which can be determined by the n + l rule.

· Electrons fill orbitals according to Hund’s rule, which states that subshells with multiple orbitals (p, d, and f) fill electrons so that every orbital in a subshell gets one electron before any of them gets a second.

o Paramagnetic materials have unpaired electrons that align with magnetic fields, attracting the material to a magnet.

o Diamagnetic materials have all paired electrons, which cannot easily be realigned; they are repelled by magnets.

· Valence electrons are those electrons in the outermost shell available for interaction (bonding) with other atoms.

o For the representative elements (those in Groups 1, 2, and 13−18), the valence electrons are found in s- and/or p-orbitals.

o For the transition elements, the valence electrons are found in s- and either d- or f-orbitals.

o Many atoms interact with other atoms to form bonds that complete an octet in the valence shell.

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