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Calcium carbonate is one of the most interesting and versatile compounds on the planet. - PowerPoint PPT Presentation


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Calcium carbonate is one of the most interesting and versatile compounds on the planet. Roughly 4% of earth’s crust The major component of rocks such as limestone and marble The white cliffs of Dover are one of the most famous natural formations of CaCO 3 .

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slide1

Calcium carbonate is one of the most interesting and versatile compounds on the planet.

  • Roughly 4% of earth’s crust
  • The major component of rocks such as limestone and marble
  • The white cliffs of Dover are one of the most famous natural formations of CaCO3.
slide3

The properties of substances & Chemical bonds

  • The properties of substances are determined in large part by the chemical bonds (O2vs N2)
  • What determines the type of bonding in each substance?
  • The electronic structures of the atoms are the key to answering the question
  • We will examine the relationship between the electronic structures of atoms and the chemical bonds they form
slide4

8.1 Chemical bonds, Lewis symbols, and the octet rule

  • Three general types of chemical bonds
  • Ionic bond: electrostatic forces
  • Covalent bond: the sharing of electrons
  • Metallic bond: free electrons
slide5

8.1 Chemical bonds, Lewis symbols, and the octet rule

  • Lewis Symbols
  • The valence electrons
  • The electrons involved in chemical bonding
  • In most atoms, they reside in the outermost occupied shell
  • The Lewis symbol
  • Diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule
slide6

8.1 Chemical bonds, Lewis symbols, and the octet rule

  • The Octet Rule
  • The noble gases have very stable electron arrangements (high ionization energies)
  • Atoms often gain, lose, or share electrons to achieve the same number of electrons as the noble gas
  • The octet rule
  • Atoms tend to gain, lose or share electrons until they are surrounded by eight valence electrons
  • The rule provides a useful framework to introduce concepts of bonding although there are many exceptions
slide7

8.2 Ionic Bonding

  • Consider a reaction of Na(s) and Cl2(g)
  • The reaction can be explained by the low ionization energy of Na and the high electron affinity of Cl
  • We can also explain the reaction by the octet rule
slide8

8.2 Ionic Bonding

  • Formation of sodium chloride
slide9

8.2 Ionic Bonding

  • Energetics of ionic bond formation
  • The formation of sodium chloride is very exothermic
  • The heat of formation of other ionic substances is also quite negative
  • The first ionization energy of Na(g) is 496 kJ/mol
  • The electron affinity of Cl(g) is −349 kJ/mol
  • If we consider only the electron transfer for the formation reaction, the energy change would be:

496 − 349 = 147 kJ/mol

How can we explain this?

slide10

8.2 Ionic Bonding

  • The Lattice Energy
  • The lattice energy is the energy required to completely separate a mole of a solid ionic compound into its gaseous ions
  • The potential energy of two interacting charged particles
  • The magnitude of lattice energies depends predominantly on the ionic charges
slide11

8.2 Ionic Bonding

  • The Lattice Energy
slide12

Practice Exercise

Which substance would you expect to have the greatest lattice energy, MgF2, CaF2, or ZrO2?

Sample Exercise 8.1 Magnitudes of Lattice Energies

Without consulting Table 8.2, arrange the following ionic compounds in order of increasing lattice energy: NaF, CsI, and CaO.

Answer: CsI< NaF < CaO.

Answer: ZrO2

slide13

8.2 Ionic Bonding

  • The lattice energy can not be determined directly by experiment
slide14

8.2 Ionic Bonding

  • Electron configurations of ions
  • The formation of Na2+ and Cl2− are energetically very unfavorable
  • Groups 1A, 2A, and 3A atoms form 1+, 2+, and 3+ ions, respectively
  • Groups 5A, 6A, and 7A atoms form 1−, 2−, and 3− ions, respectively
  • Transition metals do not observe the octet rule: Fe2+ and Fe3+
slide15

8.3 Covalent Bonding

  • A chemical bond formed by sharing a pair of electrons
  • The attractions and repulsions among electrons and nuclei in the H2 molecule
  • Quantum mechanical calculation tells us that the concentration of electron density between the nuclei leads to a net attractive force that constitutes the covalent bond holding the molecule together
slide16

8.3 Covalent Bonding

  • Lewis Structures
  • The formation of covalent bonds can be represented using Lewis symbols
  • For the nonmetals, the number of valence electrons in a neutral atom is the same as the group number

the number of

covalent bonds

1 2 3 4

slide17

8.3 Covalent Bonding

  • Multiple Bonds
  • The Lewis structures of CO2 and N2
  • Bond length: the distance between the nuclei of atoms involved in a bond
slide18

8.4 Bond Polarity and Electronegativity

  • Compare the electron-sharing of Cl2 and H2 with the sharing of H2O and NaCl
  • Polar covalent bond and nonpolar covalent bond
  • If the difference of the covalent bond-forming atoms in relative ability to attract electrons is large enough, an ionic bond is formed
  • Electronegativity is defined as the ability of an atom in a molecule to attract electrons to itself
slide19

8.4 Bond Polarity and Electronegativity

  • Electronegativity
  • Linus Pauling (1901-1994) developed the first and most widely used electronegativity scale
  • Electronegativity can be used to estimate whether a given bond will be nonpolar or polar covalent, or ionic
  • The electronegativity of an atom in a molecule is related to its ionization energy and electron affinity
  • Fluorine, the most electronegative element, has an electronegativity of 4.0
  • The least electronegative element, cesium, has an electronegativity of 0.7
slide21

8.4 Bond Polarity and Electronegativity

  • Electronegativity and bond polarity
  • The greater the difference in electronegativity between two atoms, the more polar their bond
slide22

8.4 Bond Polarity and Electronegativity

  • Electronegativity and bond polarity
slide23

8.4 Bond Polarity and Electronegativity

  • Dipole Moments
  • We can indicate the polarity of the HF molecule in two ways:
  • Polarity helps determine many of the properties of substances such as hydrogen bonding and solvation
  • How can we quantify the polarity of molecule? For the two equal and opposite charges, the dipole moment is:
  • , dipole moment (C·m or debye, D)
  • Q, charge (C)
  • r, distance (m)
  • 1 D = 3.34 × 10-30C·m
slide24

Solution

Sample Exercise 8.5 Dipole Moments of Diatomic Molecules

The bond length in the HCl molecule is 1.27 Å. (a) Calculate the dipole moment, in debyes, that would result if the charges on the H and Cl atoms were 1+ and 1–, respectively. (b) The experimentally measured dipole moment of HCl(g) is 1.08 D. What magnitude of charge, in units of e, on the H and Cl atoms would lead to this dipole moment?

slide25

8.4 Bond Polarity and Electronegativity

  • Dipole Moments
  • The actual charges on the atoms decrease from 0.41 in HF to 0.057 in HI
slide26

8.4 Bond Polarity and Electronegativity

  • Differentiating ionic and covalent bonding
  • The ability to quickly categorize the predominant bonding interactions in a substance as covalent or ionic imparts considerable insight into the properties of that substance
  • By considering the interaction between a metal and a nonmetal
  • SnCl4: colorless liquid, mp -33 ˚C, bp 114 ˚C, polar covalent
  • To use the difference in electronegativity
  • SnCl4: 1.2; NaCl: 2.1
  • MnO: 2.0, ionic; Mn2O7: 2.0, polar covalent
  • The increase in the oxidation state of a metal leads to an increase in the degree of covalent character in the bonding
slide27
Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.

If it is an anion, add one electron for each negative charge.

If it is a cation, subtract one electron for each positive charge.

PCl3

8.5 Drawing Lewis Structures

  • Lewis structures can help us understand the bonding in many compounds and are frequently used when discussing the properties of molecules

5 + 3(7) = 26

slide28

8.5 Drawing Lewis Structures

  • Arrange atoms. The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.

Keep track of the electrons: 26 - 6 = 20

slide29

8.5 Drawing Lewis Structures

  • Complete the octets around all the atoms bonded to the central atom. Hydrogen is an exception

Keep track of the electrons:

26 - 6 = 20; 20 - 18 = 2

slide30

8.5 Drawing Lewis Structures

  • Place any leftover electrons on the central atom

Keep track of the electrons:

26 - 6 = 20; 20 - 18 = 2; 2 - 2 = 0

slide31

8.5 Drawing Lewis Structures

  • If there are not enough electrons to give the central atom an octet, try multiple bonds
slide32

Sample Exercise 8.8 Lewis Structure for a Polyatomic Ion

Draw the Lewis structure for the BrO3– ion.

Solution

The total number of valence electrons is, therefore, 7 + (3  6) + 1 = 26.

For oxyanions— BrO3–, SO42–, NO3–, CO32–, and so forth—the oxygen atoms surround the central nonmetal atom.

Practice Exercise

Draw the Lewis structure for (a) ClO2–, (b) PO43–

Answers:(a)(b)

slide33

8.5 Drawing Lewis Structure

  • Formal Charge
  • The charge the atom would have if all the atoms in the molecule had the same electronegativity
  • How to assign formal charges.
    • For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms.
    • Subtract that from the number of valence electrons for that atom: the difference is its formal charge.
slide34

8.5 Drawing Lewis Structure

  • Formal Charge
  • The concept of formal charge helps us choose a preferred Lewis structure
  • How to choose the correct structure.
    • Choose the Lewis structure in which the atoms bear formal charges closest to zero
    • Choose the Lewis structure in which any negative charges reside on the more electronegative atoms

preferred

slide35

8.5 Drawing Lewis Structure

  • Formal Charge

preferred

slide36

8.5 Drawing Lewis Structure

Oxidation number

Formal charge

Actual partial charge

+1 −1

0 0

+0.18 −0.18

+0.178 −0.178

Dipole moment

slide37

8.6 Resonance Structures

  • Consider the Lewis structure of ozone, O3
  • Both Lewis structures are not corresponding to the structure of ozone
  • Resonance structure: An alternate way of drawing a Lewis dot structure for a compound
slide38

8.6 Resonance Structures

  • How can we understand the resonance structure?
  • The rules for drawing Lewis structures do not allow us to have a single structure that adequately represents the ozone molecule
slide39

8.6 Resonance Structures

Draw the resonance structures for HCO2−and NO3− .

slide40

8.6 Resonance Structures

  • Resonance in Benzene

All six C-C bonds are of equal length 1.40 Å

C-C single bond 1.54 Å; C=C double bond 1.34 Å

slide41

8.7 Exceptions to the Octet Rule

  • Odd number of elections
  • In a few molecules and polyatomic ions, such as ClO2, NO, NO2,and O2−, the number of valence electrons is odd
  • Complete pairing of these electrons is impossible
  • An octet around each atom cannot be achieved
  • NO contains 5 + 6 =11 valence electrons
slide42

8.7 Exceptions to the Octet Rule

  • Less than an octet of valence electrons
  • Consider boron trifluoride, BF3
  • In the Lewis structure of BF3, there are only six electrons around the boron atom and the formal charges are zero

Most important

  • If you try to complete the octet around boron:
slide43

8.7 Exceptions to the Octet Rule

  • More than an octet of valence electrons
  • Consider phosphorus pentachloride, PCl5
  • Elements from the third period and beyond have unfilled nd orbitals that can be used in bonding

The orbital diagram for the valence shell of a phosphorus atom

  • The larger the central atom is, the larger the number of atoms that can surround it
  • Draw the Lewis structure of SF4, ICl4−, and PO43−
slide44

8.8 Strength of Covalent Bonds

  • The stability of a molecule is related to the strengths of the covalent bonds it contains
  • The strength of a bond is measured by determining how much energy is required to break the bond
  • The bond enthalpy for a Cl-Cl bond, D(Cl-Cl), is measured to be 242 kJ/mol
slide46

8.8 Strength of Covalent Bonds

  • The bond enthalpy is always positive as energy is always required to break chemical bonds
  • A molecule with strong chemical bonds generally has less tendency to undergo chemical change than does with weak bonds
slide47

8.8 Strength of Covalent Bonds

  • Bond enthalpies and the enthalpies of reaction
  • We can use average bond enthalpies to estimate the enthalpies of reactions in which bonds are broken and new bonds are formed
  • Consider a gas-phase reaction:

CH4(g) + Cl2(g) → CH3Cl(g) + HCl(g)

slide48

8.8 Strength of Covalent Bonds

  • Bond enthalpies and the enthalpies of reaction
slide49

8.8 Strength of Covalent Bonds

  • Bond enthalpy and bond length
slide50
Enormous amounts of energy can be stored in chemical bonds → can be used as an explosives
  • Characteristics of an explosive
    • Very exothermic decomposition
    • Gaseous products
    • Rapid decomposition
    • Controllably stable
slide52

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2−

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