This chapter discusses chemical bonds and mixtures. It introduces electron-dot structures to show valence electrons and how they are involved in bonding. Ionic bonds form when ions with opposite charges are attracted to each other. Covalent bonds form when atoms share electrons. Polar covalent bonds result when electrons are shared unevenly. Molecular polarity arises if the polar bonds in a molecule do not cancel out. Most materials are mixtures that can be separated into pure substances. Solutions are homogeneous mixtures where one substance dissolves evenly throughout another. Concentration, molarity, and solubility are measures used to describe solutions.

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.

10. The Formation of Ions

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.

22. Covalent Bonds

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."

27. Polar Covalent Bonds

28. Molecular Polarity
• If the polar bonds in a molecule are facing in equal
and opposite directions, the polarity may cancel out.

29. Molecular Polarity
…Or not!

30. Molecular Polarity

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

41. The Chemist's Classification of Matter

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.

44. The Chemist's Classification of Matter

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.

49. Sucrose, C12H22O11
Solutions
= 342 g/mole

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.