Unformatted text preview: Chapter 2
Atomic Structure and Interatomic Bonding 2.2 Fundemental Concepts
Atoms are composed of positively charged protons, neutrons, and negatively charged electrons. A complete atom is electrically neutral, having the same number of protons as electrons. Atomic Number (Z) Each chemical element is characterized by the number of protons in the nucleus. Atomic Mass The sum of the masses of the protons and neutrons within the nucleus of a given atom. Isotopes The number of protons which an element may have is fixed, but the number of neutrons can vary. These elemental variations of neutrons are called isotopes. Atomic weight Corresponds to the weighted average of the atomic masses of the atom's naturally occurring isotopes. Atomic Mass Unit (amu) is 1/12 of the mass of the most common isotope of carbon, carbon 12, which has 6 protons and 6 neutrons. Mole Avogadro found that one mole of a substance contains 6.023 x 1023 atoms or molecules. 1 amu/atom = 1g/mol Ex. Iron (Fe) has an atomic weight of 55.85 amu/atom, and thus one mole of iron would weigh 55.85 grams, and would contain 6.023 x 1023 atoms. Atomic Models
The Bohr atomic model states that electrons are particles and orbit the nucleus like planets around the sun in discrete orbits. This model revolutionized our understanding of the atom and started early quantum mechanics. Atomic Models cont.
The wave-mechanical model later was developed as a refinement of the Bohr model, stating that electrons exhibit both particle-like and wave-like characteristics. Thus, Bohr's discrete orbits were further resolved into sub-levels, or quantum levels, and electrons were found to form more of a cloud around the atom based on probability instead of just following a discrete path. Electron Shells and Subshells
Electrons appear to only be able to travel in specific shells, or quantum numbers, around the nucleus. 4s Energy 3d 3p 3s 2p 2s 1s Quantum number Those shells are separated by regions which require energy to move across, thereby decreasing the probability that an electron will be found there. Within the shells there exist subshells which are also divided by areas which require energy to move across. Subshells
The s-subshell is the lowest energy subshell and can only hold two electrons. The p-subshell has the next lowest energy and can hold up to 6 electrons. The d-subshell holds up to 10 electrons. The f-subshell has the highest energy and can hold up to 14 electrons. When all of the subshells of an atom are full, and there are no subshells which are only partially full, the atom has a "stable electron configuration" and in general will not react chemically with other atoms. Ex. The noble gases, Helium, Neon, Argon, and Krypton. Valence Electrons
Valence electrons are the number of electrons that an atom would need in order to have all of its subshells full. Atoms will always fill the lower levels first leaving the outermost orbitals only partially filled if they don't have enough electrons. The valence will be either positive or negative depending on whether or not the atom will tend to gain or lose electrons in order to complete its subshells. The Periodic Table of Elements Interatomic Bonding
Primary Interatomic Bonding Ionic Bonding Covalent Bonding Metallic Bonding Hydrogen Bonding & Dipoles Secondary, van der Waals Bonding Ionic Bonding
Metallic elements easily give up their extra valence electrons to the non-metallic elements who are seeking to fill up their partially filled subshells in order to obtain stable electron configurations. When this happens the metallic elements have less electrons than protons, and so become positive ions, while the non-metallic elements have more electrons than protons and hence become negative ions. The two atoms then become bonded together by coulombic attractive forces between the positive and negative ions. Covalent Bonding
Stable electron configurations are obtained by atoms through sharing of electrons between two or more atoms. When an electron is shared it can count for both atoms. Ex. CH4 methane Carbon needs four electrons to be stable, while Hydrogen needs only two electrons. Through sharing, the carbon atom shares each of its electrons with a hydrogen atom, for each one it shares it gains one, and likewise with the hydrogen so that the five atoms all end up with a stable configuration. Metallic Bonding
Metallic bonding exists only in metals and their alloys. It is a type of sharing of electrons as in covalent bonding, except that the electrons are not bound to remain around their respective nulcei but rather are free to move throughout the entire metal, forming a "sea of electrons." As a result, the nuclei form a three dimensional array of positively charged ion cores in a negatively charged electron matrix. van der Waals Bonding
This is a secondary type of bonding which is present in all compounds to a certain degree, but they are weak compared to the primary bonding types. They exist because of dipoles, which are basically regions of positive and negative which exist around atoms or especially molecules. van der Waals Bonding cont.
Dipoles can be either permanent or induced, and result when a positive and negative region of an atom or molecule exists or is temporarily formed. These charged areas of the atom or molecule will be attracted by coulombic forces to other oppositely charged areas of other atoms or molecules forming bonds. Hydrogen Bonding
This is one of the most common, and can be one of the strongest types of van der Waals bonding. Unlike other elements which can give up electrons, but still maintain some electrons around the nucleus in a stable configuration, when hydrogen gives up its electron, only the positively charged nucleus remains. Thus, compounds of hydrogen generally form permanent dipoles, or at the very least dipoles are very easily induced. Homework
HW #1 2.3, 2.7, 2.10, 2.22 - - Due 6/14 Next Class
Read Chapter 3: The Structure of Crystalline Solids ...
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This note was uploaded on 05/03/2010 for the course ME 250-750 taught by Professor Signer during the Summer '10 term at Wichita State.
- Summer '10