chemistry unit 5
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Chemistry Unit 5. Chemical Bonding. Why Do Atoms Bond?. To become more stable like the noble gases. Octet Rule – atoms tend to gain, lose or share electrons in order to acquire a full set of valence electrons. (usually 8). Three Main Types of Bonds. - PowerPoint PPT PresentationTRANSCRIPT
Chemistry Unit 5
Chemical Bonding
Why Do Atoms Bond?
• To become more stable• like the noble gases.
• Octet Rule – atoms tend to gain, lose or share electrons in order to acquire a full set of valence electrons. (usually 8)
Three Main Types of Bonds• Ionic Bond – Atoms transfer electrons to
fill their valence shells, oppositely charged ions are formed, opposites attract.• Occurs between a metal and a nonmetal
• Covalent Bond – Atoms share electrons to fill their valence shells.• Occurs between nonmetals
• Metallic Bonds – Atoms share a “sea of electrons.” • Occurs between metal atoms
Properties of Metallic, Molecular and Ionic Compounds
Metallic Bonds
Molecular Compounds
Ionic Compounds
Insouluble
Some sol. in water/some in nonpolar
sol’n
Soluble in water
Conductor Non-Conductor
Conductive in sol’n
Mod. High Melting Points
Low Melting points
Very High Meting Points
Ionic Bonding• Ion – a charged particle
• A neutral atom becomes an ion when it loses or gains an electron.
• If an atom loses an electron, it becomes a (+) ion called a cation.
• If an atom gains an electron, it becomes a (-) ion called an anion.
Ionic Bonding
• Example
Na Cl
To become more stable, sodium must lose one electron
To become more stable, chlorine must gain one electron
Ionic Bonding
• Example
Na Cl
Sodium loses an electron and becomes an Na+1 ion.
Chlorine gains an electron and becomes a Cl-1 ion.
Opposites attract, and an ionic compound is formed…
NaCl
Try Another Example
Aluminum will become more stable if it gets rid of three electrons.
Bromine will become more stable if it receives one electron.
AlBr
Are both atoms more stable as a result of this transefer? No, Al must donate two more… where?
Aluminum & Bromine
AlBr
BrBr
Now, each atom has a full valence shell… all are more stable.
Aluminum and Bromine
AlBr
BrBr
Aluminum donated 3 e-, so it becomes Al+3
Each bromine accepted 1 e-, so they each become Br-1
The compound that forms is AlBr3
Let’s Wrap it Up• Ionic bonds are held together by electrostatic
forces.• The result of an ionic bond is called an ionic
compound.• Ionic bonds form between a metal and a nonmetal
atom due to large differences in electronegativity. (1.7 or greater)
• The nonmetal’s EN is so much greater than the metal’s EN that it removes the metal’s valence electron. Electrons are transferred.
For Example: Na and Cl
EN of Na = 0.9 EN of Cl = 3.0
Why does Sodium and Oxygen form an ionic bond?
3.0 EN of O
- 0.9 EN of Na
2.1 Difference in EN
• Difference in electronegativity is 2.1(>1.7)
• An ionic bond will form.
• Chlorine has a greater electronegativity, and is able to yank electrons away from sodium.
Covalent Bonding
O O
Each atom of Oxygen needs two more electrons to become more stable.
They will share two pairs of electrons.
A diatomic molecule of oxygen is formed.
O2
Try another example
O
H
H
Oxygen needs two electrons to become more stable.
Each atom of hyd rogen needs one more electron to become more stable.
All atoms become more stable (have full valence shells). A molecule of water is made.
H2O
Let’s Wrap it Up… Again!
• Covalent bonds are held together by a mutual need for the shared electrons (electronegativity) Their orbits overlap. Each electron is attracted to the positive charge of the opposite nucleus.
• The result of a covalent bond is called a molecule.
• Covalent bonds form between two nonmetals due to a small (or no) difference in EN. (less than 1.7)
• Neither atom’s EN is strong enough to remove the other atom’s electrons. Electrons are shared.
Polar and Nonpolar Covalent Bonds
• If one nonmetal has a greater EN than the other, it can “hog” the shared electrons. This forms a POLAR covalent bond. (EN difference greater than 0, but less than 1.7)
• If the nonmetals have the same EN, they will share equally and form NON POLAR covalent bond. (0 diff. in EN)
For Example: N and O
EN of N = 3.0 EN of O = 3.5
Why does Nitrogen and Oxygen form a Covalent Bond?
3.5 EN of Oxygen
- 3.0 EN of Nitrogen
0.5 = difference in EN
Difference in EN is less than 1.7, therefore a covalent bond will form.
Difference in EN is greater than 0, therefore the covalent bond will be polar. (Unequal sharing of e-)
One Final Example
If Chlorine bonds with Chlorine (a diatomic molecule), the difference in EN would be “0”, thus a nonpolar covalent bond will form. (Equal sharing of e-)
Molecular Geometry
• Linear molecules: atoms are connected in a straight line.• All molecules with only 2 atoms are linear.
• Many molecules with 3 atoms are also linear.
• Ex. O2, HCl, CO2
O C O
Molecular Geometry
• Bent: bonded atoms have a bent shape due to unshared pairs of electrons.• Unshared electron pairs exert a greater
repulsion force than the electron pairs in the bonds.
• Ex. H2O, NH3
H H
O
Molecular Geometry
• Tetrahedral: one atom bonded to four other atoms.• The angle between any two bonds is 109.5o.
• Ex. CH4 (methane)
C
H
H
HH
Writing Ionic Formulas
Calcium Chloride
• Locate the metal on the periodic table and write the element symbol with its oxidation number.
Ca +2
Writing Ionic Formulas
• Locate the nonmetal on the periodic table and write the element’s symbol with its oxidation number.
Cl-1
Ca+2 Cl-1
• Find the common factor between the two oxidation numbers.
• In this case, 2.• Decide how many of each ion is needed to make
the charge equal to the common factor.• In this case, 1 calcium ion (+2) and 2 chlorine ions
(-1 and –1 = -2). Compounds are neutral.• Use this number of ions as the subscript for the
element, and write the formula.• In this case, Ca Cl2.
Writing Ionic Formulas Part 2
Aluminum Oxide
• Locate the metal on the periodic table and write the element symbol with its oxidation number.
Al +3
Writing Ionic Formulas Part 2
• Locate the nonmetal on the periodic table and write the element’s symbol with its oxidation number.
O-2
Al+3O-2
• Find the common factor between the two oxidation numbers.
• In this case, 6.• Decide how many of each ion is needed to make
the charge equal to the common factor.• In this case, 2 aluminum ions (+3 and +3 = +6) and
3 oxygen ions (-2 and -2 and -2 = -6). Compounds are neutral.
• Use this number of ions as the subscript for the element, and write the formula.
• In this case, Al2O3.
Try these examples on your own.
• Sodium and Oxygen
• Lithium and Sulfur
• Aluminum and Chlorine
• Potassium and Nitrogen
• Magnesium and Fluorine
Naming Ionic Compounds
• Write the name of the metal.
• Write the name of the nonmetal with the ending changed to –ide.
Example:
Nitrogen = nitride Sulfur = sulfide
Oxygen = oxide Chlorine = chloride
Phosphorus = phosphide Iodine = iodide
Fluorine = fluoride Bromine = bromide
Naming Ionic Compounds
Al2S3
• Write the name of the metal.Aluminum
• Write the name of the nonmetal, changing the ending to –ide.
Sulfide• Name the compound.
Aluminum Sulfide
Naming Ionic Compounds
BaCl2
• Write the name of the metal.
Barium
• Write the name of the nonmetal, changing the ending to –ide.
Chloride
• Name the compound.
Barium Chloride
Try these examples on your own.
• BeF
• Li20
• B2S3
• Mg3N2
• CaCl2
Transition MetalsWtg. Formulas / Nmg. Compounds
• Most transition metals can form ions with more than one charge.
• Examples: Copper atoms can become
Cu +1 and Cu +2 ions
Iron atoms can become
Fe +2 and Fe +3 ions• Therefore, the oxidation number for the metal will be
given to you as a roman numeral in the name of the compound.
Writing Formulas w/Transition Metals
Iron (III) Oxide
• Write the symbol for the transition metal.
Ex. Fe
• Take the number in parentheses and write it as the oxidation number.
Ex. Fe +3
Writing Formulas w/Transition Metals
Iron (III) Oxide• Write the symbol for the nonmetal.
Ex. O• Look up its oxidation number on the
periodic table, and add it to the symbol.Ex. O -2
Writing Formulas w/Transition Metals
Fe +3 O –2
• Find the common factor between the two oxidation numbers. In this case = 6
• Decide how many of each ion is needed to make the charge equal to the common factor. In this case 2 Fe and 3 O ions.
• Use this number of ions as the subscript for the element, and write the formula.
Fe2O3
Copper (I) Sulfide
• Write the symbol for the transition metal.
Ex. Cu
• Take the number in parentheses, and write it as the oxidation number.
Ex. Cu +1
Copper (I) Sulfide
• Write the symbol for the nonmetal.
Ex. S
• Look up its oxidation number on the periodic table, and add it to the symbol.
Ex. S -2
Copper (I) Sulfide
Cu +1 S –2
• Find the common factor between the two oxidation numbers. In this case = 2
• Decide how many of each ion is needed to make the charge equal to the common factor. In this case 2 Cu and 1 S ion.
• Use this number of ions as the subscript for the element, and write the formula.
Cu2S
Naming Compounds w/Transition Metals
FeO• Look up the nonmetal on the periodic table.
Oxygen O-2
• Look up the metal on your ion chart. Find the possible oxidation numbers.
Fe +2 or Fe +3
Fe +2 or Fe +3 O-2
• Decide which ion will form in the proper ratio with the known charge on the oxygen ion.
FeO• Iron bonds in a 1 to 1 ratio with oxygen, therefore,
the iron ion must have a +2 charge. (Fe+2)• Name the compound, indicating the oxidation
number of the metal in parenthesis.
Iron (II) Oxide
Fe2O3
• Look up the nonmetal on the periodic table. Find its oxidation number.
Oxygen O-2
• Look up the metal on your ion chart. Find the possible oxidation numbers.
Fe +2 or Fe +3
Fe2O3
• Decide which ion will form in the proper ratio with the known charge on the oxygen ion.
Fe +2 or Fe +3 • Iron bonds in a 2 to 3 ratio with oxygen. Three
oxygen atoms will have a charge of -6. Therefore, two iron ions must equal +6. It must be Fe +3.
• Name the compound, indicating the oxidation number of the metal in parenthesis.
Iron (III) Oxide
Polyatomic IonsWriting Formulas / Naming Compounds
• A polyatomic ion is a covalent molecule that has an ionic charge. (As opposed to being a neutral molecule.)
• Poly = many• Atomic = atoms• Ion = charged particle• A charged particle that consists of more
than one atom.
Polyatomic Ions
Examples:
Sulfide Sulfate = SO4-2
Nitride Nitrate = NO3-1
Phosphide Phosphate = PO4-3
Chloride Chlorate = ClO3-1
• Notice the ending has changed to –ate.
Polyatomic Ions
Examples:
Sulfide Sulfite = SO3-2
Nitride Nitrite = NO2-1
Phosphide Phosphite = PO3-3
Chloride Chlorite = ClO2-1
• Notice the ending has changed to –ite.
Polyatomic ions
• Not all polyaomic ions end in -ate or -ite.
• Some other examples:
Ammonium NH4+1
Hydroxide OH-1
• Some Polyatomic ions contain more than two elements. Ex. Acetate = C2H3O2
-1
Calcium Phosphite
• Write the symbol for the metal. Add the oxidation number from the periodic table.
Ca+2
• Write the formula for the polyatomic ion from the ion chart. Add its oxidation number.
PO3-3
Calcium PhosphiteCa +2 PO3 -3
• Determine the common factor of the two oxidation numbers. In this case, 6.
• Decide how many of each ion is needed to equal the common factor. In this case, 3 calcium ions and 2 phosphate ions.
• Write these numbers as the subscript for each ion.
Ca3(PO3)2
• Notice that the polyatomic ion must be placed in parenthesis or, instead of 2 phosphate ions, you would have 32 Oxygen atoms and 1 Phosphorus atom.
Ca3PO32
An Additional Example
• Aluminum HydroxideAl+3 OH-1
• The least common factor is 3. Therefore, 1 aluminum ion will bond with 3 hydroxide ions to form a neutral compound.
Al(OH)3
• If you omitted the parenthesis, you would not have 3 hydroxide ions. Instead you would have 3 hydrogen atoms and one oxygen atom. AlOH3
Try these!
• Write formulas for the following compounds. Lithium sulfate Calcium acetate Aluminum nitrite Magnesium phosphate Sodium carbonate
Answers
• Li2SO4 Lithium sulfate
• Ca(C2H3O2)2 Calcium acetate
• Al(NO2)3 Aluminum nitrite
• Mg3(PO4)2 Magnesium phosphate
• Na2CO3 Sodium carbonate
Naming Compoundsw/ Polyatomic Ions
KClO3
• Write the name of the metal.Potassium
• Write the name of the polyatomic ion from the ion chart.
Chlorate• Name the compound.
Potassium Chlorate
Mg3(SO3)2
• Name the metal.
Magnesium
• Name the polyatomic ion from the ion chart.
Sulfite
• Name the compound.
Magnesium Sulfite
Try these
• Write names for the following compounds.
CaCO3
Al2(SO3)3
Ca(ClO2)2
K3PO4
Mg(OH)2
Answers
• CaCO3 Calcium carbonate
• Al2(SO3)3 Aluminum sulfite
• Ca(ClO3)2 Calcium chlorate
• K3PO4 Potassium phosphate
• Mg(OH)2 Magnesium hydroxide
Covalent MoleculesUse prefixes to designate the number of atoms of each element used in the molecule.
Prefix Number
mono 1
di 2
tri 3
tetra 4
penta 5
Prefix Number
hexa 6
hepta 7
octa 8
nona 9
deca 10
Writing Covalent Formulas
Dinitrogen Pentoxide
• Write the symbol of each element.
N O• Add the subscript as indicated by the prefixes.
N2O5
Writing Covalent Formulas
Carbon Dioxide
• Write the name of each element.
C O• Add the subscripts as indicated by the prefixes.
CO2
Try these examples on your own.
• Sulfur dioxide
• Sulfur monoxide
• Carbon tetrachloride
• Dihydrogen dioxide
• Nitrogen triiodide
Naming Covalent Molecules
NH3
• Write the name of the first nonmetal using its subscript as a prefix.
Nitrogen (No prefix written for the first element IF it is a one.)
• Write the name of the second nonmetal using its subscript as a prefix and change the ending to -ide.
Trihydride• Name of the molecule: Nitrogen trihydride
Naming Covalent Molecules Pt.2
P2O5
• Write the name of the first nonmetal using the subscript as a prefix.
Diphosphorus• Write the name of the second nonmetal using the
subscript as a prefix and change the ending to -ide.
Pentoxide• Name the molecule: Diphosphorus Pentoxide
Try a few examples on your own.
• CO
• CO2
• SF2
• PI3
• H2O
Naming Acids
• Acids are water solutions of certain hydrogen compounds.
• There are two main types of acids:• Binary Acids – Hydrogen + a nonmetal
• Tertiary Acids – Hydrogen + a polyatomic ion
Binary Acids
• To name the acid.
Hydro _____ ic acid. (The blank is the root of the nonmetal.)
For example:
HCl = Hydrochloric acid
HI = Hydroiodic acid
HBr = Hydrobromic acid
Binary Acids
• To write the formula:
Hydrofluoric acid
Write the symbol for hydrogen.
Write the symbol for the nonmetal.
Use the ox. numbers to figure out the ratio.
H +1 F-1
HF
Binary Acids
• Hydrosulfuric acidWrite the symbol for hydrogen.
Write the symbol for the nonmetal.
Use the ox. Numbers to figure out the ratio.
H+1 S-2
H2S
Tertiary Acids
Hydrogen plus a polyatomic ion.
If the anion ends in –ate, _____ic acid.
If the anion ends in –ite, _____ ous acid.
(I ate it and it was icky.)
(Rite ous!)
Tertiary Acids
• To name an -ate acid:
H3PO4 = Phosphoric acid (phosphate ion)
HClO3 = Chloric acid (chlorate ion)
HNO3 = Nitric acid (nitrate ion)
• To name an –ite acid:
H2SO3 = Sulfurous acid (sulfite ion)
HClO2 = Chlorous acid (chlorite ion)
HNO2 = Nitrous acid (nitrite ion)
Tertiary Acids• To write formulas for ___ ic acids:
Carbonic acid: Write symbol for hydrogen. Write polyatomic –ate ion.
H+1 CO3-2 (carbonate ion)
H2CO3
• To write formulas for ___ous acids:Clorous acid: Write the symbol for hydrogen
Write the polyatomic –ite ion.
H+1 ClO2-1 (chlorite ion)
HClO2
Properties of Water
• The unique properties of water are due to the strong intermolecular HYDROGEN BONDS that are formed between the polar water molecules. (Opposite poles attract. The positive hydrogen end of one molecule of water attracts to the negative oxygen end of another molecule of water.)
Hydrogen Bonding
• This hydrogen bonding is more extensive in ice than it is in liquid water. For this reason, ice is less dense than water.
• Hydrogen bonding in water gives it a high surface tension.
Properties of Water• Water is called the universal solvent due to
its polarity. (Pulls ionic compounds apart)• Water has a very high specific heat. (It
takes a lot of energy to overcome the forces between the molecules and make them move faster and heat up) This moderates climates near large bodies of water.
Properties of Water
• Water molecules are adhesive, they stick to other things. (Forms a meniscus, leaves containers wet)
• Water molecules are also cohesive, they stick to each other. (Capillary action – draws water up through plants, trees, etc.)