Properties of Covalent Compounds

11 Key excerpts on "Properties of Covalent Compounds"

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  eBook - PDF

  Comparative Inorganic Chemistry

  • Bernard Moody(Author)
  • 2013(Publication Date)
  • Arnold

    (Publisher)

  5 Bonding and the structures displayed by elements and their compounds The general physical properties of compounds related to bond type The nature of the bonding in a compound and of the geometrical pattern adopted by the ions or molecules in a solid, will largely determine the physical properties ofthat substance. While distinc-tive properties associated with ionic and covalent bonding may be discerned, there is a gradual merging of characteristics when the compounds of a large number of elements are compared. This gradual transition is not altogether unexpected. When fused or dissolved in water, an ionic com-pound will conduct electricity. The current is carried through the liquid by the ions which gain their mobility when the compound is melted or dispersed in a solvent. Unless a reaction occurs with the solvent, covalent substances yield non-conducting liquids. In the crystal lattice of an ionic compound, each ion is surrounded by oppositely charged ions, the number depending on the particular pattern adopted in the crystal. Strong electrical forces hold the ions in position although each atom oscillates by virtue of its thermal energy. Considerable energy is required to overcome the forces of attraction and ionic compounds usually melt at high temperatures and are non-volatile. On the other hand, covalent molecules, each electrically neutral, are held by much weaker intermolecular forces. Therefore, fusion, boiling and sublimation are relatively easy to accomplish. To illustrate this point, the melting-points of the fluorides formed by the elements of Period 3, sodium-sulphur, are shown in Table 5.1.

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  General Chemistry: Atoms First

  • Young, William Vining, Roberta Day, Beatrice Botch(Authors)
  • 2017(Publication Date)
  • Cengage Learning EMEA

    (Publisher)

  6 Covalent Bonding Unit Outline 6.1 Covalent Bonding and Lewis Structures 6.2 Properties of Covalent Bonds 6.3 Resonance and Bond Properties In This Unit… We will examine chemical bonding in detail in this unit and the next. Here we apply what you have learned about atomic structure, elec-tron configurations, and periodic trends to the chemical bonds formed between atoms. This unit and the next primarily address covalent bond-ing; we examined ionic bonding briefly in Ionic and Covalent Compounds (Unit 5), and we will do so in more detail in The Solid State (Unit 13). We will also examine the forces that exist between individual particles, called intermolecular forces, in Intermolecular Forces and the Liquid State (Unit 12). Vasilyev/Shutterstock.com Copyright 2018 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. WCN 02-300 Unit 6 Covalent Bonding 132 6.1 Covalent Bonding and Lewis Structures 6.1a Fundamentals of Covalent Bonding A covalent bond is characterized by the sharing of valence electrons by two adjacent atoms. This type of bonding occurs most typically between nonmetal elements such as carbon, hydrogen, oxygen, and nitrogen. For example, consider a simple covalently bonded molecule, H 2 (Interactive Figure 6.1.1). When two isolated H atoms are at a great distance from one another, they feel no attractive or repulsive forces. However, as the atoms approach more closely, the attractive and repulsive forces between the two atoms become important. At very short dis-tances, repulsive forces become more important than attractive forces, and the atoms repel. Attractive forces between the hydrogen atoms result from the interaction between the positively charged nucleus (represented by the element symbol) on one hydrogen atom and the negatively charged electron (represented by a dot) on the other hydrogen atom.

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  Introduction to Molecular Science

  • Ageetha Vanamudan(Author)
  • 2023(Publication Date)
  • Delve Publishing

    (Publisher)

  Source: By DynaBlast - Created with Inkscape, CC BY-SA 2.5, https://com- mons.wikimedia.org/w/index.php?curid=995735 Covalent Bonding 115 7.1 DO YOU UNDERSTAND WHY COVALENT BONDS FORM? When the difference in electronegativity between two atoms is insufficient for electron transit, a covalent bond is created. Electronegativity is a measure of an atom’s capacity to attract electrons. Atoms create covalent bonds to strengthen their stability by forming a whole electron shell by sharing their outermost (valence) electrons.The attraction between positively charged nuclei and shared electrons is stronger than the repulsion between them in covalent bonds. This attraction aids in the retention of the molecules. The amount of energy required to break the connection, or the amount of energy required to separate the connected atoms determines the strength of a covalent bond. Figure 7.2: Structure of the covalent bonding of methane. Source: By Original: Benjah-bmm27 Vector: Jynto - Own work based on: File:Methane-CRC-MW-dimensions-2D.png, Public Domain, https://com- mons.wikimedia.org/w/index.php?curid=12422898 7.2 COVALENT BOND CHARACTERISTICS In the structure of stable covalent compounds, covalent bonding is typically the controlling factor. A covalent bond has the following characteristics and qualities: • It is possible to make multiple covalent bonds between two atoms when composed of two or more nonmetals, or one nonmetal and a metalloid. Breaking bonds requires a lot of energy, which is why it is so tough (Krylov & Gill, 2013). Directional or Isomerism is the phenomenon in which one chemical formula may represent several distinct molecules. Introduction to Molecular Science 116 • Covalent compounds have low melting and boiling temperatures, do not conduct electricity, and are insoluble in polar solvents such as water. Covalent substances can be dissolved using nonpolar solvents such as benzene and toluene.

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  Basic Concepts of Environmental Chemistry

  • Des W. Connell(Author)
  • 2005(Publication Date)
  • CRC Press

    (Publisher)

  To understand why matter exists as solids, liquids, or gases and how they physically evaporate, dissolve, and generally distribute in the environment, we must start by considering the molecular nature of chemical compounds and the way atoms are bound to one another. First, we will look at the nature of chemical bonds, and this will give an insight into physical-chemical properties that will provide an understanding of environmental properties.

  TABLE 2.1 Properties of Some Substances in the Environment

  Substances	Occurrence in the Environment	Normal Physical State	Boiling Point (°C) (Atmospheric Pressure)	Melting Point (°C)
  Oxygen	21% of atmosphere	Gas	−183	−218
  Nitrogen	78% of atmosphere	Gas	−196	−210
  Water	Oceans, lakes, rivers	Liquid	100	0
  Common salt	3.5% of seawater	Solid	1413	801
  Quartz	Rocks, sand, geological strata	Solid	2230	1610

  The most important type of bonding is the covalent bond. Usually with this type of bonding, two atoms react together, with each contributing one electron to form a bond. Thus, two atoms of hydrogen can react to form a hydrogen molecule with one covalent bond. Thus,

  H ⋅ + ⋅ H → H − H

  Each covalent bond consists of two electrons moving rapidly between the hydrogen atoms in a defined space, as illustrated in Figure 2.2 . Each electron holds a full negative charge, and if the electrons spend equivalent times near the two atoms in the bond, this results in no difference in charge between the two ends of the bond. This can be interpreted as the electron density around the two atoms being symmetrical, leading to a nonpolar bond. However, a different situation applies with the hydrogen chloride bond. In this bond, the chlorine tends to attract electrons as shown in Figure 2.3 . This attraction is not sufficient to cause the electron to remain permanently with the chlorine atom, but causes the electron to spend more of its time when it is moving between the two atoms toward the chlorine end of the bond. This results in a small partial negative charge (denoted as δ−) occurring on the chlorine atom and leads to a small partial positive charge (denoted as δ+) occurring at the hydrogen atom. These factors are illustrated with the hydrogen chloride bond in Figure 2.3 . The outcome of this effect is that the bond becomes polar and has a dipole moment or polarity.

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  Foundations of College Chemistry

  • Morris Hein, Susan Arena, Cary Willard(Authors)
  • 2016(Publication Date)
  • Wiley

    (Publisher)

  11.4 Predicting Formulas of Ionic Compounds • Chemical compounds are always electrically neutral. • Metals lose electrons and nonmetals gain electrons to form compounds. • Stability is achieved (for representative elements) by attaining a noble gas electron configuration. 11.5 The Covalent Bond: Sharing Electrons • Covalent bonds are formed when two atoms share a pair of electrons between them: • This is the predominant type of bonding in compounds. • True molecules exist in covalent compounds. • Overlap of orbitals forms a covalent bond. • Unequal sharing of electrons results in a polar covalent bond. 11.6 Electronegativity • Electronegativity is the attractive force an atom has for shared electrons in a molecule or polyatomic ion. • Electrons spend more time closer to the more electronegative atom in a bond forming a polar bond. • The polarity of a bond is determined by the electronegativity difference between the atoms involved in the bond: • The greater the difference, the more polar the bond is. • At the extremes: • Large differences result in ionic bonds. • Tiny differences (or no difference) result(s) in a nonpolar covalent bond. KEY TERM ionization energy KEY TERM Lewis structure KEY TERM ionic bond KEY TERMS covalent bond polar covalent bond KEY TERMS electronegativity nonpolar covalent bond dipole 246 CHAPTER 11 • Chemical Bonds: The Formation of Compounds from Atoms • A molecule that is electrically asymmetrical has a dipole, resulting in charged areas within the molecule. H : Cl H Cl + δ - δ hydrogen chloride -δ +δ • If the electronegativity difference between two bonded atoms is greater than 1.7–1.9, the bond will be more ionic than covalent. • Polar bonds do not always result in polar molecules. 11.7 Lewis Structures of Compounds Problem-Solving Strategy Writing a Lewis Structure 1. Obtain the total number of valence electrons to be used in the structure by adding the number of valence electrons in all the atoms in the molecule or ion.

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  Understanding Basic Chemistry

  The Learner's Approach

  • Kim Seng Chan, Jeanne Tan(Authors)
  • 2014(Publication Date)
  • WSPC

    (Publisher)

  The physical properties of simple discrete molecular compounds and macromolecular/giant covalent compounds are totally dissimilar. Thus, we need to discuss them separately.

  Volatility of simple discrete molecular compounds

  Simple covalent compounds, such as CO2 and H2 O, have low melting and boiling points. The bonds between the molecules are weak intermolecular bonds/ intermolecular forces (e.g., van der Waals’ forces or hydrogen bond); thus, the amount of heat energy required to break them is low.

  Melting and boiling are phase changes for a substance. When a covalent compound undergoes a phase change, it still retains its molecular entity. What differs would be the extent of interactions it has with other molecules and also its degree of freedom of motion.

  The strong covalent bonds between the atoms in a molecule are not readily broken down by heat, thus the molecule remains intact when the substance changes state. This is a common misconception that students have!

  Examples:

  Methane (CH4 ) m.p. = −182°C

  Sulfur (S8 ) m.p. = 144°C

  Volatility of macromolecular compounds

  Some non-metallic substances such as diamond and silicon have high melting and boiling points, which means that the attractive forces binding the particles together are very strong. For melting to occur, a great amount of energy in the form of heat is required to overcome the strong covalent bonds between the atoms.

  In fact, diamond, an allotrope of carbon, is made up of a big network of carbon atoms each covalently bonded to four other carbon atoms. Silicon has a similar structure as diamond, whereas the macromolecular structure of silicon dioxide (SiO2 ) consists of each Si atom covalently bonded to four oxygen atoms, and each oxygen atom is in turn covalently bonded to two Si atoms. So, take note that the chemical formula for macromolecular compounds, such as SiO2 , is in fact an empirical formula and not the typical molecular formula for simple discrete molecular compounds.

  Examples:

  Silicon	m.p. = 1,650°C
  Diamond	m.p. = 3,700°C
  Graphite	m.p. = 3,300°C

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  Environmental Inorganic Chemistry for Engineers

  • James G. Speight(Author)
  • 2017(Publication Date)
  • Butterworth-Heinemann

    (Publisher)

  In the present context of inorganic chemistry, the manner in which molecular inorganic molecules and ions pack in the solid state not only is important for the formation of crystals but also contributes significantly to the conduction, magnetic, and nonlinear optical properties of these molecular inorganic compounds in the solid state. Furthermore, inorganic compounds present coordination geometries different from those found for carbon. For example, although four-coordinate carbon is nearly always tetrahedral, both tetrahedral and square planar shapes occur for four-coordinate compounds of both metals and nonmetals. When metals are the central atoms, with anions or neutral molecules bonded to them (frequently through nitrogen, oxygen, or sulfur), these are usually designated as (called) coordination complexes; when carbon is the element directly bonded to metal atoms or ions, the chemicals are organometallic compounds.

  4.3 Physical Properties

  As already noted, inorganic compounds form ionic bonds have high melting points and are made from either single elements or compounds that do not include carbon and hydrogen. In solutions, they break down into ions that conduct electricity. Organic compounds have a carbon-based structure with covalent bonding and are often volatile in nature. Even in liquid state, they do not conduct electricity unless they are salts formed with inorganic acids and bases.

  Due to the type of bonding, most inorganic compounds are polar compared with organic chemicals (Speight, 2017b ) and have a strong tendency to ionize in water or when subjected to electrolysis. The low volatility of inorganic compounds is due to the strength of the intramolecular bonding; the property of low volatility is increased in situations in which other attractive forces within (and between) the molecules add to the strength of the various bonding arrangements due to weak bonds, such as hydrogen bonds and van der Waals forces.

  Examples of properties that might be considered to be relevant to identification and characterization of inorganic compounds include melting point and/or boiling point, crystal shape, and color. Inorganic compounds tend to have high melting points, and while some inorganic compounds are solids with accessible melting points and some are liquids with reasonable boiling points; they are not the exhaustive tabulations of melting point data and boiling point data for inorganic compounds that exist for organic chemicals (Speight, 2017b

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  Plastics Materials

  • J A Brydson(Author)
  • 1999(Publication Date)
  • Butterworth-Heinemann

    (Publisher)

  5

  Relation of Structure to Chemical Properties

  5.1 INTRODUCTION

  It is sometimes stated that a given material has ‘a good chemical resistance’, or alternatively the material may be stated to be poor or excellent in this respect. Such an all-embracing statement can be little more than a rough generalisation, particularly since there are many facets to the behaviour of polymers in chemical environments.

  There are a number of properties of a polymer about which information is required before detailed statements can be made about its chemical properties. The most important of these are:

  (1) The solubility characteristics. (2) The effect of specific chemicals on molecular structure, particularly in so far as they lead to degradation and cross-linking reactions. (3) The effect of specific chemicals and environments on polymer properties at elevated temperatures. (4) The effect of high-energy irradiation. (5) The aging and weathering of the material. (6) Permeability and diffusion characteristics. (7) Toxicity.

  Before dealing with each of these aspects, it is useful to consider, very briefly, the types of bonds which hold atoms and molecules together.

  5.2 CHEMICAL BONDS

  The atoms of a molecule are held together by primary bonds. The attractive forces which act between molecules are usually referred to as secondary bonds, secondary valence forces, intermolecular forces or van der Waals forces.

  Primary bond formation takes place by various interactions between electrons in the outermost shell of two atoms resulting in the production of a more stable state. The three main basic types of primary bond are ionic, covalent and coordinate.

  An ionic bond is formed by the donation of an electron by one atom to another so that in each there is a stable number of electrons in the outermost shell (eight in the case of most atoms). An example is the reaction of sodium and chlorine (Figure 5.1

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  Introduction to General, Organic, and Biochemistry

  • Morris Hein, Scott Pattison, Susan Arena, Leo R. Best(Authors)
  • 2014(Publication Date)
  • Wiley

    (Publisher)

  11.4 PREDICTING FORMULAS OF IONIC COMPOUNDS • Chemical compounds are always electrically neutral. • Metals lose electrons and nonmetals gain electrons to form compounds. • Stability is achieved (for representative elements) by attaining a noble gas electron configuration. 11.5 THE COVALENT BOND: SHARING ELECTRONS • Covalent bonds are formed when two atoms share a pair of electrons between them: • This is the predominant type of bonding in compounds. • True molecules exist in covalent compounds. • Overlap of orbitals forms a covalent bond. • Unequal sharing of electrons results in a polar covalent bond. 11.6 ELECTRONEGATIVITY • Electronegativity is the attractive force an atom has for shared electrons in a molecule or polyatomic ion. • Electrons spend more time closer to the more electronegative atom in a bond forming a polar bond. • The polarity of a bond is determined by the electronegativity difference between the atoms involved in the bond: • The greater the difference, the more polar the bond is. • At the extremes: • Large differences result in ionic bonds. • Tiny differences (or no difference) result(s) in a nonpolar covalent bond. C H A P T E R 1 1 R E V I E W KEY TERM ionization energy KEY TERM Lewis structure KEY TERM ionic bond KEY TERMS covalent bond polar covalent bond KEY TERMS electronegativity nonpolar covalent bond dipole 240 CHAPTER 11 • Chemical Bonds: The Formation of Compounds from Atoms • A molecule that is electrically asymmetrical has a dipole, resulting in charged areas within the molecule. −δ +δ H : Cl H Cl � � � � hydrogen chloride • If the electronegativity difference between two bonded atoms is greater than 1.7–1.9, the bond will be more ionic than covalent. • Polar bonds do not always result in polar molecules. 11.7 LEWIS STRUCTURES OF COMPOUNDS PROBLEM-SOLVING STRATEGY: Writing a Lewis Structure 1.

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  Chemistry

  Structure and Dynamics

  • James N. Spencer, George M. Bodner, Lyman H. Rickard(Authors)
  • 2011(Publication Date)
  • Wiley

    (Publisher)

  and bronze), or intermolecular compounds (such as Li 3 As) are metallic bonds. Distinguishing among solids that are primarily held together by covalent, ionic, or metal bonds is useful because it allows us to predict many of the physical prop- erties of the solid. 9.2 Molecular and Network Covalent Solids MOLECULAR SOLIDS The iodine (I 2 ) that dissolves in alcohol to make the antiseptic known as tinc- ture of iodine, the cane sugar (C 12 H 22 O 11 ) found in a sugar bowl, and the poly- ethylene used to make garbage bags all have one thing in common. They are all examples of compounds that are molecular solids at room temperature. Water and bromine are liquids that form molecular solids when cooled slightly; H 2 O freezes at 0°C and Br 2 freezes at 7°C. Many substances that are gases at room temperature will form molecular solids when cooled far enough; F 2 , at the extreme right of the bond-type triangle in Figure 9.1, freezes to form a molecu- lar solid at 220°C. Molecular solids contain both intramolecular bonds and intermolecular forces, as described in Chapter 8. The atoms within the individual molecules are held together by relatively strong intramolecular covalent bonds. Molecular solids are therefore found in the covalent region of a bond-type triangle. The molecules in these solids are held together by much weaker intermolecular forces. Because intermolecular forces are relatively weak, molecular solids are often soft sub- stances with low melting points. Dry ice, or solid carbon dioxide, is a perfect example of a molecular solid. The van der Waals forces holding the CO 2 molecules together are weak enough that at dry ice sublimes at a temperature of 78ºC––it goes directly from the solid to the gas phase. Changes in the strength of the van der Waals forces that hold molecular solids together can have important consequences for the properties of the solid.

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  Visualizing Everyday Chemistry

  • Douglas P. Heller, Carl H. Snyder(Authors)
  • 2015(Publication Date)
  • Wiley

    (Publisher)

  single bond One pair of shared electrons serving as a covalent bond between two atoms. double bond Two pairs of shared electrons serving as two covalent bonds between two atoms. triple bond Three pairs of shared electrons serving as three covalent bonds between two atoms. + ) ) ) ) ) ) + 2 2 2 2 2 + 1 1 1 1 1 1 2 a. Each atom shares one valence electron to create a single bond. b. Each atom shares two valence electrons to create a double bond. c. Each atom shares three valence electrons to create a triple bond. Ionic and Covalent Bonding 65 vaporize until it reaches an astonishing 1413°C (2575°F) (Figure 3.14 on the next page). The major reason for these different responses to changes in temperature lies in the different attractive forces present in ionic and covalent compounds. In crys- talline table salt, the sodium cations and the chloride an- ions are closely packed in an orderly arrangement that produces the crystal lattice of the salt we use with our food. In this sodium chloride lattice, the positively charged sodium cations and the negatively charged chlo- ride anions lie close to each other in a repeating pattern, with each attraction. When we heat both salt crystals and ice, we can see the effects of the powerful ionic forces that keep an ionic crystal intact as well as the effects of much weaker intermolecular forces of a covalent substance. Well below room temperature, H 2 O, a covalent mol- ecule, exists as ice, the frozen form of water. As the tem- perature rises to 0°C (32°F), solid ice melts to liquid water, which remains a liquid (at normal atmospheric pressure) until it reaches a temperature of 100°C (212°F), where it boils and converts to steam, or water vapor. Table salt, an ionic compound (NaCl), requires far higher temperatures for similar transformations.

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