1. The Formation of Ions
An ion is an atom that has lost or gained one or more electrons.
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2. The Formation of Ions
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3. H H O H+ O H H H The Formation of Ions
A molecular ion is typically formed by the loss or gain of a hydrogen ion, H+. H H O + H+ O H H H
Water
Hydrogen ion
Hydronium ion, H3O+
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4. Ionic Bonds
An ionic bond is the electrical force of attraction between oppositely charged ions.
Na+ F-
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5. Ionic Bonding Used to show ration of ions.
Uses element symbols to indicate atoms present in the compound and subscripts to indicate ratios of atoms.
Example: Aluminum oxide  Al2O3
2 aluminum ions for every 3 oxide ions.
Example: Lithium oxide  Li2O
2 lithium ions for every oxide ion.
If no subscript 1 is assumed.
Not all compounds found in 1:1 ration.
Ratio depends on the charges of the ions.
The compound has to be neutral.

6. Ionic Bonding Example:
Calcium forms ions with:
2+ charge
Fluorine forms ions with:
1- charge
You need two fluorine atoms for every one calcium atom to get a neutral compound.
CaF2

7. Ionic Bonding Cation (+) written on the left, anion (-) on the right.
Ratio needs to make the compound neutral
Use “Criss-cross” method
Example: Lithium Oxide= Li1+ O2-  Li2O

8. Ionic Bonding
If possible the formula should be simplified to show the smallest whole-number ratio of the ions:
Mg2+ + O2- Mg2O2  MgO
For compounds with polyatomic ion, the polyatomic ion should be placed in parenthesis:
Ca2+ + NO3-  Ca(NO3)2

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10. Sample Problems Write the formulas for: Sodium Bromide
Magnesium Phosphate
Potassium Iodide
Barium Chloride

11. Ionic Bonds CHECK YOUR NEIGHBOR
What is the chemical formula for a compound made of aluminum ions, Al3+, and oxygen ions, O2–?
AlO
Al3O2
Al2O3
Al6O6
Explain your answer to your neighbor.
C. Al2O3
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12. Ionic Bonds CHECK YOUR NEIGHBOR
What is the chemical formula for a compound made of magnesium ions, Mg2+, and oxygen ions, O2–?
MgO
Mg2O2
Mg4O4
any of the above
Explain your answer to your neighbor.
A. MgO
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13. 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.
So what affects one atom affects all the atoms bonded together.
This mobility of electrons accounts for many metallic properties.
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14. Metallic Bonds
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15. Metallic Bonds
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16. Metallic Bonds An alloy is a mixture of metallic elements.
77% Cu, 12% Zn, 7% Mn, 4% Ni
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17. Covalent Bonds
A covalent bond is the type of electrical attraction in which atoms are held together by their mutual attraction for shared electrons.
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18. Covalent Bonds
Electrons are shared not transferred.
Covalent bond – formed by a shared pair of electrons between two atoms.
Found in paper, plastic, plant matter, human tissue…..

19. Describing Covalent Bonds
Described by the octet rule
Example:
A Fluorine atoms has 7 valence electrons
Each fluorine atom in molecular fluorine shares 1 electron with the other fluorine atom – results in a covalent bond.
Ammonia (NH3)

20. Single covalent bonds – single bonds
Multiple Bonds
Single covalent bonds – single bonds
2 atoms share exactly one pair of electrons
Double covalent bonds – double bond
Two pairs of shared electrons
Triple covalent bond – triple bond
Three pairs of shared electrons

21. Covalent Bonds
There are two electrons in a single covalent bond.
The covalent bond is represented using a straight line.
F F
F — F
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22. Covalent Bonds
The number of covalent bonds an atom can form equals its number of unpaired valence electrons.
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23. Covalent Bonds
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24. Covalent Bonds Multiple covalent bonds are possible.
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25. Double Bonds Example: Formaldehyde (H2CO)
Double bonds – 2 pairs of shared electrons between atoms.

26. Triple Bonds Example: Acetylene C2H2
Triple bond – three pairs of shared electron pairs between atoms.
Shown with 6 dots between element symbols.

27. Dots or Dashes Another way to represent Lewis Structure: Example:
A single line represents a single bond
Double line = double bond….
Lines represent bonds, dots represent unshared pairs of electrons.
Example:
Ammonia NH3
Formaldehyde

28. Properties of Covalent Bonds
A covalent bond is a shared pair of electrons between 2 atoms.
Electrons are not always shared equally
Different atoms have different electronegativity
Measure of an atom’s attraction for electrons.
Shared electrons are more strongly attracted to the atom that is more electronegative.

29. Properties of Covalent Bonds
When one atom is significantly more electronegative than another the covalent bond is polar.
The atom with greater electronegativity gains a slightly negative charge because it has a slight excess of electron density
The shared electrons spend more time next to the more electronegative atom.

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31. Properties of Covalent Bonds
Example: water molecule
Oxygen is more electronegative than hydrogen.
Oxygen atom in a water molecule has a slight negative charge
Hydrogen atom has a slight positive charge.

32. Polarity
To show the partial charge the Greek symbol δ (delta) is used:
δ+ partial positive, δ- partial negative
Arrows are also used to show the charge difference.

33. Polarity
A bond between 2 atoms that have similar electronegativities will be nonpolar.
Have relatively weak attractions to other nonpolar molecules
Nonpolar covalent bond
F-F

34. Polarity
Electonegativities can be used to predict if a bond will be polar, nonpolar, or ionic.
Electronegativity difference between two atoms < 0.4 = nonpolar
Electronegativity between 0.4 and 2.0 = polar
Electronegativity > 2.0 = ionic
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36. Polar Covalent Bonds
Electrons in a covalent bond are shared evenly when the two atoms are the same.
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37. Polar Covalent Bonds
Electrons in a covalent bond may be shared unevenly, however, when the bonded atoms are different.
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38. Polar Covalent Bonds
Electronegativity is the ability of a bonded atom to pull on shared electrons. Greater electronegativity means greater "pulling power."
High
Low
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40. Polar Covalent Bonds
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41. Polar Covalent Bonds CHECK YOUR NEIGHBOR
Which is heavier: carbon dioxide, CO2, or water, H2O?
Carbon dioxide is heavier.
Water is heavier.
Both have the same number of atoms, so they weigh the same.
It depends on other factors.
Explain your answer to your neighbor.
A. Carbon dioxide is heavier.
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42. Molecular Polarity
If the polar bonds in a molecule are facing in equal and opposite directions, the polarity may cancel out.
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44. Molecular Polarity
…Or not!
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45. Molecular Polarity Because water is bent the bonds don’t cancel out.
It is a polar molecule
So it has a “strong” attraction to other polar molecules
This effects its physical and chemical properties
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46. Molecular Polarity
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47. Molecular Polarity CHECK YOUR NEIGHBOR
Water has such a relatively high boiling point because water
is such a heavy substance.
is transparent so that heat passes right through it.
contains three atoms per molecule.
molecules are so sticky.
Explain your answer to your neighbor.
D. molecules are so sticky.
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48. Molecular Attractions
An ion dipole attraction is the attraction between an ion and a dipole.
Example: NaCl in water H H H O O O H H H
Na+ H
Cl- H O O H H H O H
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49. Molecular Attractions
An dipole-dipole attraction is the attraction between two dipoles.
Example: cohesive forces within water
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50. Dipoles
A dipole can also attract a molecule that is ordinarily not a dipole.
When a dipole approaches a nonpolar molecule, its partial charge either attracts or repels the electrons of the other particle.

51. Molecular Attractions
A dipole–induced dipole attraction is the attraction between a dipole and an induced dipole.
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52. Molecular Attractions CHECK YOUR NEIGHBOR
Is it possible for a fish to drown?
no, because fish breathe water
yes, when the water contains too little oxygen
no, because water is 88.8% oxygen by mass
yes, when the water is not moving
Explain your answer to your neighbor.
B. yes, when the water contains too little oxygen
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53. Molecular Attractions CHECK YOUR NEIGHBOR
A nonpolar material, such as oxygen, O2, is soluble in a polar material, such as water, H2O, by way of
ion–dipole attractions.
dipole–dipole attractions.
dipole–induced dipole attractions.
all of the above
Explain your answer to your neighbor.
C. dipole–induced dipole attractions.
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54. Molecular Attractions
A fourth molecular attraction is the induced dipole–induced dipole, which occurs between nonpolar molecules. (Weakest attraction)
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55. Molecular Attractions
Nonpolar atoms are attracted to each other by these "momentary" dipoles.
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56. Molecular Attractions
The larger the atom, the stronger the "momentary" dipole.
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57. Molecular Attractions
The tiny nonpolar fluorine atoms in Teflon provide very weak
attractions, which is why Teflon provides a "nonstick" surface.
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58. Molecular Attractions
Small nonpolar molecules are easier to pull apart then large nonpolar molecules
Like small pieces of velcro vs large pieces of velcro
Why do you need soap to clean your clothes?
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59. Molecular Attractions
So, how do the gecko's sticky feet stay so clean?
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60. Molecular Attractions CHECK YOUR NEIGHBOR
Which type of molecular attraction takes the least amount of energy to break apart?
ion–dipole attractions
dipole–dipole attractions
dipole–induced dipole attractions
induced dipole–induced dipole attractions
Explain your answer to your neighbor.
D. induced dipole-induced dipole attractions
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61. 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.
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62. Most Materials Are Mixtures
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63. The Chemist's Classification of Matter
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64. 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.
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65. 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.
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66. The Chemist's Classification of Matter
Heterogeneous mixture is a mixture that is composed of more than one phase, and are unevenly mixed.
The different phases are separated by definite boundaries called interfaces
Example: granite, milk

67. Two or more substances that can be easily separated by physical means form a mixture.
Heterogeneous mixture—mixture of different and easily distinguishable materials

68. The Chemist's Classification of Matter
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69. The Chemist's Classification of Matter
Homogeneous mixtures:
Solution: all components in the same phase
Suspension: different components in different phases
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70. Homogeneous mixture consist of only one phase, and are evenly mixed.
Example: sugar, salt, seawater, quartz, air
Homogeneous mixture—contains two or more gaseous, liquid, or solid substances blended evenly; also called a solution

71. The Chemist's Classification of Matter CHECK YOUR NEIGHBOR
Is the air in your house a homogeneous or a heterogeneous mixture?
homogeneous, because it is mixed very well
heterogeneous, because of the dust particles it contains
homogeneous, because it is all at the same temperature
heterogeneous, because it consists of different types of molecules
Explain your answer to your neighbor.
B. heterogeneous, because of the dust particles it contains.
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72. The Chemist's Classification of Matter CHECK YOUR ANSWER
Is the air in your house a homogeneous or a heterogeneous mixture?
homogeneous, because it is mixed very well
heterogeneous, because of the dust particles it contains
homogeneous, because it is all at the same temperature
heterogeneous, because it consists of different types of molecules
B. heterogeneous, because of the dust particles it contains.
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73. Solutions
A solution is a homogenous mixture consisting of ions or molecules.
A solution consists of a solute (dissolved material) in a solvent (dissolving material)
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.
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74. Solutions Solutions are not necessarily always liquid
Both air and glass are solutions

75. 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.
Amount of solute
Concentration =
Amount of solution
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76. The factor-label method treats labels as factors.
These factors divide out.
There are two main rules to solving science problems with the factor-label method:
1. Always carry along your units with any measurement you use.
2. You need to form the appropriate labeled ratios (equalities).

77. SIGNIFICANT DIGITS
In addition and subtraction, the answer may contain only as many decimal places as the measurement having the least number of decimal places.
In multiplication and division, the answer may contain only as many significant digits as the measurement with the least number of significant digits.

78. Significant Digit Rules
1) ALL non-zero numbers (1,2,3,4,5,6,7,8,9) are ALWAYS significant.
2) ALL zeroes between non-zero numbers are ALWAYS significant.
3) ALL zeroes which are SIMULTANEOUSLY to the right of the decimal point AND at the end of the number are ALWAYS significant.
4) ALL zeroes which are to the left of a written decimal point and are in a number >= 10 are ALWAYS significant.

79. Example Problem:
How many centimeters in 2 meters?
You will see from the metric conversion chart that 1 meter = 100 cm
we turn this into a ratio by writing it like this:

80. Or

81. Once you have the equalities you must pick the one that will cancel out the units leaving the desired units.
Then multiply your starting quantity (2 meters) by the equality that will give you your desired units.

82. As a rule of thumb your problem set up should look like this:

83. Practice Problems
1. How many wheels on 350 Ford pickups (use the equality 1 pickup = 4 tires) - the starting units are pickups, the ending units need to be wheels.

84. Practice Problems
2. How many millimeters in 34 hectometers (use the equality 10,000 mm = 1 hectometer)?
Sometimes you will need to multiply by more than one ratio to get to your desired units, you can do this by using linking units. Your setup will look like this:

85. 3. How many inches in 1 meter given the equality 1 inch = 2
3. How many inches in 1 meter given the equality 1 inch = 2.54 cm and 1 meter = 100 cm? (note the linking unit in this problem is cm)

86. 4. How many grams in 150 pounds given the equalities 1 pound = 0
4. How many grams in 150 pounds given the equalities 1 pound = kg and 1 kg = 1000 grams?

87. Solutions
A mole is a super-large number, 6.02  1023, used to measure numbers of atoms or molecules, also called Avogadro's number.
Substance Formula Mass
The formula mass of a substance expressed in grams contains 1 mole.
Carbon, C 12
Oxygen, O2 32
Carbon dioxide, CO2 44
Sucrose, C12H22O11 342
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88. Solutions Sucrose, C12H22O11 = 342 g/mole
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89. Solutions CHECK YOUR NEIGHBOR
Water, H2O, has a formula mass of 18. How many moles of water are there in 18 grams of water?
0.5 mole
1 mole
9 moles
18 moles
Explain your answer to your neighbor.
B. 1 mole
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90. Solutions CHECK YOUR NEIGHBOR
How many grams of water, H2O, are there in 2 moles of water?
1 gram
9 grams
18 grams
36 grams
Explain your answer to your neighbor.
D. 36 grams
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91. Molarity (M) – the most common expression of solution concentration.
Defined as the number of moles of solute dissolved in each liter of solution.

92. Molarity
Example:
What is the molarity of a solution formed by mixing 10.0 g H2SO4 with enough water to make 100.0mL of solution?
- convert ml to L
- complete the molarity equation
- convert grams to moles

93. Practice Problems
What is the molarity of the solution produced when 145 g of sodium chloride is dissolved in sufficient water to prepare 2.75 L of solution?

94. Solutions
ppm is a unit of concentration expressed in milligrams of solute in per liters of solution.
1 part solute
1 milligram solute
1 ppm = =
1,000,000 parts solution
1 liter solution
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95. Solubility
Solubility is the ability of a solute to dissolve in a solvent.
A solute that has appreciable solubility is said to be soluble.
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96. Solubility A precipitate is solute that comes out of solution.
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97. Temperature
Solutions of gases in liquids are greatly affected by changes in temperature.
As temperature increases, the kinetic energy of the solute gas becomes greater.
As the temperature increases, the solubility of a gas in a liquid decreases.

98. Temperature
Solubility of solids in liquids increases as the temperature increases.
Relationship depends on the energy change during solution formation.
If the temperature drops when the solute and solvent are mixed, raising the temperature will increase the solubility.

99. Solubility CHECK YOUR NEIGHBOR
The amount of oxygen, O2, dissolved in the waters of the Arctic Ocean is –––––––– the amount of oxygen dissolved in warm tropical waters.
greater than
about equal to
less than
It depends on other factors.
Explain your answer to your neighbor.
A. greater than
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100. Chapter 12 Review CHECK YOUR NEIGHBOR
By mass, water is 88.88% oxygen. So, why can't we breathe water?
Hint: What is the elemental formula for the oxygen we breathe and the chemical formula for water?
Explain your answer to your neighbor.
C. discover the property called inertia.
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