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Chapter 12:
Chemical Bonds
and Mixtures
Principles of Science II
This lecture will help you understand:
• Electron-Dot Structures
• The Formation of Ions
• Ionic Bonds
• Metallic Bonds
• Covalent Bonds
• Polar Covalent Bonds
• Molecular Polarity
• Molecular Attractions
• Most Materials Are Mixtures
• The Chemist's Classification of Matter
• Solutions
• Solubility
Electron-Dot Structures
• Atoms bond together through their electrons. To
learn about bonding, therefore, we need to know
something about how the electrons in an atom are
organized.
• Electrons behave as though they are contained
within seven concentric shells.
Electron-Dot Structures
• The numbers indicate the maximum number of
electrons each shell may contain.
Note:
• This is a "conceptual model"
and not a representation of
what an atom "looks like."
• Rather, it helps us to
understand how the
electrons in atoms behave.
Electron-Dot Structures
• The shells are more easily
drawn in two dimensions.
• Each atom has its own
configuration of electrons.
Elements in the same group
have similar configurations,
which is why they have
similar properties.
Electron-Dot Structures
• Valence electrons are electrons in the outermost
shell of an atom. These are the ones that can
participate in chemical bonding.
• An electron-dot structure is a notation that shows
the valence electrons surrounding the atomic
symbol.
Kr
Electron-Dot Structures
• Special Note
– For heavier atoms, some valence electrons are
more available than others. Krypton, for example,
has 18 valence electrons, but only eight of these
are typically shown in an electron-dot structure.
These are the eight that extend farthest away
from the nucleus.
Electron-Dot Structures
• Note that elements in the same group have the
same electron-dot structure.
The Formation of Ions
• An ion is an atom that has lost or gained one or
more electrons.
The Formation of Ions
H
O
H
H+
Water Hydrogen ion
H
O
H
H
+
Hydronium ion, H3O+
The Formation of Ions
• A molecular ion is typically formed by the loss or
gain of a hydrogen ion, H+
.
Na+ F-
Ionic Bonds
• An ionic bond is the electrical force of attraction
between oppositely charged ions.
Chapter 12 chemical bonds and mixtures
Chapter 12 chemical bonds and mixtures
Ionic Bond Formation
1) An electrically neutral sodium atom loses its valence
electron to an electrically neutral chlorine atom.
2) This electron transfer results in two oppositely charged ions
3) The ions are held together by an ionic bond.
Chapter 12 chemical bonds and mixtures
Metallic Bonds
• Outer electrons in metal atoms are held only weakly
by the nucleus.
• This weak attraction allows the electrons to move
about quite freely.
• This mobility of electrons accounts for many metallic
properties.
Metallic Bonds
• An alloy is a mixture of metallic elements.
Covalent Bonds
A covalent bond is the type of electrical attraction in
which atoms are held together by their mutual
attraction for shared electrons.
F — FF F
Covalent Bonds
• There are two electrons in a single covalent bond.
• The covalent bond is represented using a straight
line.
Covalent Bonds
• The number of covalent bonds an atom can form
equals its number of unpaired valence electrons.
Covalent Bonds
Covalent Bonds
• Multiple covalent bonds are possible.
Polar Covalent Bonds
• Electrons in a covalent bond are shared evenly
when the two atoms are the same.
Polar Covalent Bonds
• Electrons in a covalent bond may be shared
unevenly, however, when the bonded atoms are
different.
High
Low
Polar Covalent Bonds
• Electronegativity is the ability of a bonded atom to
pull on shared electrons. Greater electronegativity
means greater "pulling power."
Polar Covalent Bonds
Molecular Polarity
• If the polar bonds in a molecule are facing in equal
and opposite directions, the polarity may cancel out.
Molecular Polarity
…Or not!
Molecular Polarity
Na+
Cl-
H
O
H
H
O
H
H
O
H
O
H
H
O
H
H H
H
O
Molecular Attractions
• An ion dipole attraction is the attraction between an
ion and a dipole.
– Example: NaCl in water
Molecular Attractions
• An dipole-dipole attraction is the attraction between
two dipoles.
– Example: cohesive forces within water
Molecular Attractions
• A dipole–induced dipole attraction is the attraction
between a dipole and an induced dipole.
Molecular Attractions
• A fourth molecular attraction is the induced dipole–
induced dipole, which occurs between nonpolar
molecules.
Nonpolar atoms are attracted to each other by these "momentary" dipoles.
Molecular Attractions
The larger the atom, the stronger the "momentary" dipole.
Molecular Attractions
Molecular Attractions
The tiny nonpolar fluorine atoms in Teflon provide very weak
attractions, which is why Teflon provides a "nonstick" surface.
So, how do the gecko's sticky feet stay so clean?
Molecular Attractions
Most Materials Are Mixtures
• A pure substance is a material that consists of only
one type of element or compound.
• A mixture is a collection of two or more pure
substances.
• It can be separated by physical means.
Most Materials Are Mixtures
The Chemist's Classification of Matter
The Chemist's Classification of Matter
• Pure materials consist of a single element or
compound.
• Impure materials consist of two or more elements or
compounds.
• Mixtures may be heterogeneous or homogeneous.
The Chemist's Classification of Matter
• In heterogeneous mixtures, the different
components can be seen as individual substances.
• In homogeneous mixtures, the composition is the
same throughout.
The Chemist's Classification of Matter
The Chemist's Classification of Matter
• Homogeneous mixtures:
– Solution: all components in the same phase
– Suspension: different components in different
phases
© 2013 Pearson Education, Inc.
Solutions
• A solution is a homogenous mixture consisting of
ions or molecules.
• A solvent is the major component of a solution.
• A solute is the minor component of a solution.
• If a solution is saturated, then no more solute will
dissolve in it.
Concentration
Amount of solute
Amount of solution
Solutions
• Concentration is a measure of the amount of solute
dissolved in solution.
• A solution with more solute than solution is called
concentrated.
• A solution with more solution than solute is called
dilute.
=
The formula mass of a
substance expressed in
grams contains 1 mole.
Substance Formula Mass
Carbon, C 12
Oxygen, O2 32
Carbon dioxide, CO2 44
Sucrose, C12H22O11 342
Solutions
• A mole is a super-large number, 6.02 × 1023
, used to
measure numbers of atoms or molecules, also
called Avogadro's number.
Sucrose, C12H22O11
Solutions
= 342 g/mole
Molarity
moles of Solute
liter of Solution
Solutions
• Molarity is a unit of concentration expressed in
moles of solute per liter of solution.
=
1 ppm
1 part solute
1,000,000 parts solution
1 milligram solute
1 liter solution
=
Solutions
• ppm is a unit of concentration expressed in
milligrams of solute in per liters of solution.
=
Solubility
• Solubility is the ability of a solute to dissolve in a
solvent.
• A solute that has appreciable solubility is said to be
soluble.
Solubility
• A precipitate is solute that comes out of solution.

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Chapter 12 chemical bonds and mixtures

  • 1. Chapter 12: Chemical Bonds and Mixtures Principles of Science II
  • 2. This lecture will help you understand: • Electron-Dot Structures • The Formation of Ions • Ionic Bonds • Metallic Bonds • Covalent Bonds • Polar Covalent Bonds • Molecular Polarity • Molecular Attractions • Most Materials Are Mixtures • The Chemist's Classification of Matter • Solutions • Solubility
  • 3. Electron-Dot Structures • Atoms bond together through their electrons. To learn about bonding, therefore, we need to know something about how the electrons in an atom are organized. • Electrons behave as though they are contained within seven concentric shells.
  • 4. Electron-Dot Structures • The numbers indicate the maximum number of electrons each shell may contain. Note: • This is a "conceptual model" and not a representation of what an atom "looks like." • Rather, it helps us to understand how the electrons in atoms behave.
  • 5. Electron-Dot Structures • The shells are more easily drawn in two dimensions. • Each atom has its own configuration of electrons. Elements in the same group have similar configurations, which is why they have similar properties.
  • 6. Electron-Dot Structures • Valence electrons are electrons in the outermost shell of an atom. These are the ones that can participate in chemical bonding. • An electron-dot structure is a notation that shows the valence electrons surrounding the atomic symbol.
  • 7. Kr Electron-Dot Structures • Special Note – For heavier atoms, some valence electrons are more available than others. Krypton, for example, has 18 valence electrons, but only eight of these are typically shown in an electron-dot structure. These are the eight that extend farthest away from the nucleus.
  • 8. Electron-Dot Structures • Note that elements in the same group have the same electron-dot structure.
  • 9. The Formation of Ions • An ion is an atom that has lost or gained one or more electrons.
  • 11. H O H H+ Water Hydrogen ion H O H H + Hydronium ion, H3O+ The Formation of Ions • A molecular ion is typically formed by the loss or gain of a hydrogen ion, H+ .
  • 12. Na+ F- Ionic Bonds • An ionic bond is the electrical force of attraction between oppositely charged ions.
  • 15. Ionic Bond Formation 1) An electrically neutral sodium atom loses its valence electron to an electrically neutral chlorine atom. 2) This electron transfer results in two oppositely charged ions 3) The ions are held together by an ionic bond.
  • 17. Metallic Bonds • Outer electrons in metal atoms are held only weakly by the nucleus. • This weak attraction allows the electrons to move about quite freely. • This mobility of electrons accounts for many metallic properties.
  • 18. Metallic Bonds • An alloy is a mixture of metallic elements.
  • 19. Covalent Bonds A covalent bond is the type of electrical attraction in which atoms are held together by their mutual attraction for shared electrons.
  • 20. F — FF F Covalent Bonds • There are two electrons in a single covalent bond. • The covalent bond is represented using a straight line.
  • 21. Covalent Bonds • The number of covalent bonds an atom can form equals its number of unpaired valence electrons.
  • 23. Covalent Bonds • Multiple covalent bonds are possible.
  • 24. Polar Covalent Bonds • Electrons in a covalent bond are shared evenly when the two atoms are the same.
  • 25. Polar Covalent Bonds • Electrons in a covalent bond may be shared unevenly, however, when the bonded atoms are different.
  • 26. High Low Polar Covalent Bonds • Electronegativity is the ability of a bonded atom to pull on shared electrons. Greater electronegativity means greater "pulling power."
  • 28. Molecular Polarity • If the polar bonds in a molecule are facing in equal and opposite directions, the polarity may cancel out.
  • 31. Na+ Cl- H O H H O H H O H O H H O H H H H O Molecular Attractions • An ion dipole attraction is the attraction between an ion and a dipole. – Example: NaCl in water
  • 32. Molecular Attractions • An dipole-dipole attraction is the attraction between two dipoles. – Example: cohesive forces within water
  • 33. Molecular Attractions • A dipole–induced dipole attraction is the attraction between a dipole and an induced dipole.
  • 34. Molecular Attractions • A fourth molecular attraction is the induced dipole– induced dipole, which occurs between nonpolar molecules.
  • 35. Nonpolar atoms are attracted to each other by these "momentary" dipoles. Molecular Attractions
  • 36. The larger the atom, the stronger the "momentary" dipole. Molecular Attractions
  • 37. Molecular Attractions The tiny nonpolar fluorine atoms in Teflon provide very weak attractions, which is why Teflon provides a "nonstick" surface.
  • 38. So, how do the gecko's sticky feet stay so clean? Molecular Attractions
  • 39. Most Materials Are Mixtures • A pure substance is a material that consists of only one type of element or compound. • A mixture is a collection of two or more pure substances. • It can be separated by physical means.
  • 40. Most Materials Are Mixtures
  • 42. The Chemist's Classification of Matter • Pure materials consist of a single element or compound. • Impure materials consist of two or more elements or compounds. • Mixtures may be heterogeneous or homogeneous.
  • 43. The Chemist's Classification of Matter • In heterogeneous mixtures, the different components can be seen as individual substances. • In homogeneous mixtures, the composition is the same throughout.
  • 45. The Chemist's Classification of Matter • Homogeneous mixtures: – Solution: all components in the same phase – Suspension: different components in different phases
  • 46. © 2013 Pearson Education, Inc. Solutions • A solution is a homogenous mixture consisting of ions or molecules. • A solvent is the major component of a solution. • A solute is the minor component of a solution. • If a solution is saturated, then no more solute will dissolve in it.
  • 47. Concentration Amount of solute Amount of solution Solutions • Concentration is a measure of the amount of solute dissolved in solution. • A solution with more solute than solution is called concentrated. • A solution with more solution than solute is called dilute. =
  • 48. The formula mass of a substance expressed in grams contains 1 mole. Substance Formula Mass Carbon, C 12 Oxygen, O2 32 Carbon dioxide, CO2 44 Sucrose, C12H22O11 342 Solutions • A mole is a super-large number, 6.02 × 1023 , used to measure numbers of atoms or molecules, also called Avogadro's number.
  • 50. Molarity moles of Solute liter of Solution Solutions • Molarity is a unit of concentration expressed in moles of solute per liter of solution. =
  • 51. 1 ppm 1 part solute 1,000,000 parts solution 1 milligram solute 1 liter solution = Solutions • ppm is a unit of concentration expressed in milligrams of solute in per liters of solution. =
  • 52. Solubility • Solubility is the ability of a solute to dissolve in a solvent. • A solute that has appreciable solubility is said to be soluble.
  • 53. Solubility • A precipitate is solute that comes out of solution.

Editor's Notes

  1. Figure 12.6
  2. Table 12.1
  3. Figure 12.7
  4. Figure 12.9