8 Key excerpts on "Bonding"

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  eBook - ePub

  AP® Chemistry All Access Book + Online + Mobile

  • Kevin Reel, Derrick C. Wood, Scott A. Best(Authors)
  • 2014(Publication Date)
  • Research & Education Association

    (Publisher)

  6

  Chemical Bonding

  Intramolecular Forces: Bonds Between Atoms

  Bonds are the forces of attraction that hold atoms together. There are many types of Bonding including ionic, metallic, and covalent bonds. You can figure out the difference between the Bonding types if you look at what role the valence electrons are playing in the chemical bond—because Bonding is all about the valence electrons. Many of the electrons in an atom have no impact on Bonding because they are located close to the nucleus, and thus are called core electrons . In general, the valence electrons are the outermost s-shell and p-shell electrons in an electron configuration. For transition metals, the outermost d-shell electrons will also play a role. Elements will typically form bonds in order to have eight electrons in the valence shell, which is called the octet rule . The vast majority of chemical bonds that occur obey the octet rule, although a significant number of exceptions to the octet rule exist; these exceptions will be covered later in this chapter.

  Ionic Bonds

      •   Ionic Bonding is a bond between a cation and an anion held together by electrostatic attractions. Coulomb’s law dictates that oppositely charged particles are attracted to one another, and this is the fundamental principle behind ionic Bonding.

      •   In an ionic bond, an electron is removed from the least electronegative atom to form a positively charged ion (cation). This electron is then transferred to a more electronegative atom to form a negatively charged ion (anion).

      •   Ionic bonds form in order to fulfill the octet rule for the elements involved with the bond. The metal loses electrons to have a filled shell. The nonmetal gains electrons to have a filled shell.

     

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  The Science For Conservators Series

  Volume 1: An Introduction to Materials

  • The Conservation Unit Museums and Galleries Commission(Author)
  • 2008(Publication Date)
  • Routledge

    (Publisher)

  The electrons can thus move freely from one atom to another because none of the overlapping outer orbitals is full. In other words, the outer electrons belong to all the atoms. The mobile electrons act as a cohesive force preventing the positive metal ions from pushing each other apart. Just as with the ionic bond there is no distinct group of atoms which can be identified as a molecule. This type of structure can extend indefinitely in any direction. Often the most stable structure (and hence the most likely one) will be that in which there is the greatest overlap of the orbitals of one atom with those of its neighbours. This is achieved in many metals (such as copper, silver, and gold) by a regular pattern called close-packing. The regular array of repeating units in three dimensions suggests that metals are crystalline. It is unusual to see individual metal crystals in isolation, but solid metal objects are made up of large numbers of small crystals joined together. The boundaries between these crystals can be seen under a microscope.

  Figure 4.18 Photomicrograph, showing the individual crystals in a sample of brass – 60% copper, 40% zinc.

  Three types of bond?

  Covalent, ionic and metallic bonds have been described as if they formed three distinct categories. We have seen that really these three are just definable points in a wide range of Bonding behaviour.

  Covalent Bonding describes not only the equal sharing of an electron pair between two atoms but may involve electron sharing over larger numbers of atoms (as in SO2 ). The extreme form of this is the completely delocalised electron structure of metals.

  In a covalent bond the electrons may not be shared equally between the atoms if one is more electronegative than another (oxygen in H2 O). The extreme form of unequal sharing is when there is complete electron transfer from one atom to another, as is found in the ionic bond.

  D Physical properties related to Bonding

  The materials of which objects are made have different physical characteristics which distinguish them from each other. Materials are often chosen for a particular job because of these distinguishing features. Copper, which conducts electricity well, is used in wires to carry current, while plastics such as polyvinyl chloride, which do not conduct electricity, are used as insulation for the copper wire. Solvents with low boiling points are used for dry-cleaning because they will evaporate rapidly from the textile, once the cleaning has been finished.

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  eBook - ePub

  Understanding General Chemistry

  • Atef Korchef(Author)
  • 2022(Publication Date)
  • CRC Press

    (Publisher)

  8 Chemical Bonding and Molecular Geometry

  DOI: 10.1201/9781003257059-8

  8.1 Objectives

  At the end of the present chapter, the student will be able to:

  1. Recognize the different types of chemical Bonding.
  2. Draw the Lewis structure of different atoms and ions.
  3. Recognize the different molecular geometries.
  4. Determine the geometry of molecules.

  8.2 Chemical Bonding

  Chemical Bonding (intramolecular forces) is defined as the attractive forces that hold two or more chemical constituents (atoms and ions) together in different chemical species. The electrons involved in Bonding are usually those in the outermost (valence) shell.

  There are three types of chemical bonds:

  1. Covalent bonds result from sharing electrons between the atoms. The covalent bond is usually found between non-metals. Depending on the electronegativities of the atoms involved in the bond, the covalent bond can be polar or non-polar (Figure 8.1 ). Polar covalent bonds occur when the difference in electronegativity between the two constituent atoms is between 0.4 and 2.  Examples of polar covalent bonds: The bond between hydrogen and chlorine in HCl is a polar covalent bond. Chlorine (Cl) is more electronegative than hydrogen (H). The shared pair of electrons between Cl and H is attracted toward the chlorine atom. Because of this, the hydrogen atom has a partially positive charge (δ+ ) and the chlorine atom has a partially negative charge (δ− ). In addition, the bond between hydrogen and bromine in HBr, the bond between hydrogen and fluorine in HF and the bond between hydrogen and oxygen in water H2 O are polar covalent bonds.  Examples of non-polar covalent bonds: The bond between two iodine atoms in I2 is a non-polar covalent bond. The bonds between hydrogen (H) and carbon (C) in CH4 are non-polar covalent bonds.  Note that the non-polar covalent bond is found in homonuclear molecules such as Br2 , Cl2 , O2 , I2 and in hydrocarbons (Cn H2n + 2

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  General Chemistry for Engineers

  • Jeffrey Gaffney, Nancy Marley(Authors)
  • 2017(Publication Date)
  • Elsevier

    (Publisher)

  Chapter 3

  Chemical Bonding—The Formation of Materials

  Abstract

  This chapter covers chemical Bonding between atoms and ions and how this affects the chemical properties of the elements. Which elements form ions and the typical charges on the ions are explained. Ionic Bonding and covalent Bonding are compared in terms of the octet rule and valence bond theory. Polar and nonpolar covalent bonds are explained and their relationship to both electron group geometry and molecular geometry is stressed. Polyatomic ions are described as a mixed ionic, covalent species. Molecular orbital theory is introduced to explain magnetism, bond order, and hybridization, which will be important in later discussions of the chemistry of carbon. Intermolecular forces, including hydrogen Bonding, are discussed with a special Case Study focusing on the special properties of water.

  Keywords

  Ionic Bonding; Covalent Bonding; Octet rule; Polyatomic ions; Dipole moment; Molecular orbitals; Hybridization; Resonance; Molecular geometry; Intermolecular forces

  Outline

  3.1  

  Atoms and Ions

  3.2  

  Ionic Bonding

  3.3  

  Covalent Bonding

  3.4  

  Mixed Covalent/Ionic Bonding

  3.5  

  Molecular Orbitals

  3.6  

  Molecular Geometry

  3.7  

  Molecular Polarity

  3.8  

  Intermolecular Forces

  Important Terms

  Study Questions

  Problems

  3.1 Atoms and Ions

  A neutral atom that loses one or more electrons becomes a positively charged ion. This positively charged ion is known as a cation (from the Greek word katá , meaning “down”). A neutral atom that gains one or more electrons has a negative charge and is known as an anion (from the Greek word ánō , meaning “up”). The number of electrons an element will gain or lose is also a periodic property and can generally be predicted from its position in the periodic table as shown in Fig. 3.1 . Atoms will gain or lose electrons to form ions that have electronic configurations which are more stable than the electronic configurations of the parent atoms. For most elements, this means that they will either gain or lose the number of electrons needed to achieve a closed valence shell. Remember from Table 2.9 of Chapter 2

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  BIOS Instant Notes in Physical Chemistry

  • Gavin Whittaker, Andy Mount, Matthew Heal(Authors)
  • 2000(Publication Date)
  • Taylor & Francis

    (Publisher)

  Valence theories attempt to describe the number, nature and strength of chemical bonds between atoms. It also describes the geometric arrangement of the bonds, and so the shapes of molecules. The more sophisticated valence theories yield information about the electrical, magnetic, and spectroscopic properties of molecules.

  Elementary valence theories invoke two principal bond types. In ionic Bonding , electrostatic interactions generate bonds between ions formed by electron transfer from one element to the other. In covalent Bonding two elements are held together by shared electrons in order that both may adopt an energetically favorable electron configuration. In reality, both are extreme forms of the same Bonding phenomenon. Pure covalent bonds are formed by elements with identical electronegativities, with more ionic Bonding character being introduced to the bond as the electronegativity difference between the elements increases (see Topic H4 ). Even in extreme cases of ionic Bonding, the degree of covalent character may still be quite high.

  Two complementary theories were originally developed to explain the number and nature of covalent bonds (Lewis theory) and the shapes of molecules (VSEPR theory). More sophisticated theories have superseded these approaches for detailed investigations, but they remain useful in semi-empirical and non-rigorous discussions of molecular Bonding.

  Lewis theory

  The Lewis theory of covalent Bonding may be regarded as an elementary form of valence bond theory. It is nonetheless useful for describing covalent molecules with simple covalent bonds, and works successfully in describing the majority of, for example, organic compounds. Lewis theory recognizes both the free energy gains made in the formation of complete atomic electron shells, and the ability of atoms to achieve this state by sharing electrons. The sharing process is used as a description of covalent bonds.

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  Foundations for Teaching Chemistry

  Chemical Knowledge for Teaching

  • Keith S. Taber(Author)
  • 2019(Publication Date)
  • Routledge

    (Publisher)

  Often, teaching schemes begin with covalent Bonding and use the metaphor of a covalent bond being a ‘shared’ pair of electrons – sometimes without exploring what exactly sharing means in this context. A teacher cannot be surprised that students may later talk of atoms ‘stealing’ electrons and the like if a social metaphor is presented as if a satisfactory description. As suggested earlier, it would be better to focus on how the negative electrons can hold the positively charged atomic cores in the structure when there is a balance of attractive and repulsive forces. A student might reasonably ask how the two negative electrons can be considered to behave as a pair if they repel each other – a question that cannot be addressed in any detail without invoking that elephant we may prefer to pretend is not in the room.

  Covalent Bonding is often illustrated with misleading hypothetical schemes for how the molecules came about: thus, methane is often shown as being formed from an isolated carbon atom and four isolated hydrogen atoms – allowing students to acquire the mistaken notion that bonds form in order to allow atoms to follow the octet rule (chemical reaction mechanisms are discussed in Chapter 11 ). The students may be new enough to the subject not to spot this trick and so to ask where these radical atoms originated: a teacher should know better than to use such deceptive and unscientific devices in explanations. A teacher should only explain a reaction as starting with non-bonded atoms if they can show the students that reaction using reactants in that form – but I doubt any school chemical stores can supply such reagents!

  The ionic bond is based on the attractions between charges

  The traditional teaching scheme will often then move on to ionic Bonding. This is a more complex situation because it does not involve discrete links between adjacent atomic cores but rather relates to the overall regular structures of myriad cations and anions that allow forces to be balanced by placing positive ions closer to negative ions than other positive ions and vice versa. (That may seem to give an overall attractive force, not an equilibrium, but one has to consider that once ions come very close, then the outer electrons, all negative, have a small separation so contribute strong repulsions.) Often, to simplify matters, teaching models ignore the lattice originally and focus on molecule-like parts of the structure – such as one Na+ and one Cl− ion in NaCl. Whilst this is certainly simpler, the aim is to explain NaCl, which has an extensive lattice of ions, not hypothetical Na+ -Cl− ion pairs (which may exist at a low level in the vapour phase, or hydrated as a minority component in very concentrated solutions – but have no role in the solid). If we can treat water as composed of H2 O molecules to a good first approximation (e.g., ignoring the very low proportion of ions present, see Chapter 8 ), then there is no place for Na+ -Cl−

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  Chemistry

  Concepts and Problems, A Self-Teaching Guide

  • Richard Post, Chad Snyder, Clifford C. Houk(Authors)
  • 2020(Publication Date)
  • Jossey-Bass

    (Publisher)

  3 Periodic Properties and Chemical Bonding In Chapter 1 you learned that the elements in a horizontal row of the periodic table show regular variation in properties from left to right. The elements are arranged in the table in order by increasing atomic number (reading the table left to right, line by line, in the way you are reading this paragraph). The reason this arrangement works so well is that all atoms consist of electrons, protons, and neutrons. The neutrons and protons are in the nucleus with electrons arranged in “shells” around the nucleus. Why did we consider the electronic arrangement of atoms in such detail? Because the chemical properties of an element depend upon the number of electrons in its outermost shell, the energy levels of its outermost electrons, and the size of the atom. These details of atomic structure determine what kinds and how many chemical bonds can be formed by an atom. In this chapter we discuss several properties not mentioned in Chapter 1 that depend upon the outermost shell electronic structure. We will review electron configuration and introduce new “dot” symbols. We will then discuss whether atoms gain, lose, or share electrons, and how many, when they combine to form new substances. The major portion of the chapter is devoted to the types of chemical bonds (ionic, covalent, polar covalent) formed between atoms in chemical compounds

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  Nursing HESI A2

  a QuickStudy Laminated Reference & Study Guide

  • April Michelle Davis(Author)
  • 2018(Publication Date)
  • QuickStudy Reference Guides

    (Publisher)

  Hydrogen Bonding: Hydrogen bonded to electro­negative atoms
• Van der Waals forces: The sum of small force inter­actions between molecules not in covalent, ionic, or hydrogen bonds
• London dispersion force: Temporary dipoles created by the normal movement of electrons