Unit 3: Chemical Formulas

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• Unit 3 Chemical Formulas
• and
• Bonding

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Electron Dot diagrams

• Electron dot diagrams show the valence electrons
  around an atom. In most molecules and compounds
  a complete octet is achieved for each atom
• Most monatomic ions have an electron
  configuration of noble gases

Al
N
7 valence e-s
e- ?
8 valence e-s
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Drawing Lewis Dot Structures

• To visualize valence e-, we will use Lewis Dot
  Structures.
• Step 1 The element symbol represents the nucleus
  and all e- except valence.
• Step 2 From the periodic table, determine the
  number of valence e-.
• Step 3 Each side of symbol represents an
  orbital. Draw two dots on one side, then one for
  each of the remaining three sides. Additional
  electrons should then be paired.

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Lewis Dot Structures

• Ex carbon
• step 1 C
• step 2 4 valence e-s
• step 3
• C

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Lewis Dot Structures

• Ex bromine
• step 1 Br
• step 2 7 valence e-s
• step 3
• Br

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Chemical BondingWhat holds things together?
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Lets examine the melting point of compounds
across two periods.
What is the trend?
Nonconductive
Chlorides of Period 2 Chlorides of Period 2 Chlorides of Period 2 Chlorides of Period 2 Chlorides of Period 2 Chlorides of Period 2 Chlorides of Period 2 Chlorides of Period 2
compound LiCl BeCl2 BCl3 CCl4 NCl3 OCl2 Cl2
melting point 610 415 -107 -23 -40 -121 -102
Chlorides of Period 3 Chlorides of Period 3 Chlorides of Period 3 Chlorides of Period 3 Chlorides of Period 3 Chlorides of Period 3 Chlorides of Period 3 Chlorides of Period 3
compound NaCl MgCl2 AlCl3 SiCl4 PCl3 SCl6 Cl2
melting point 801 714 193 -69 -112 -51 -102
low
high
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Bonding

• How can we explain the melting point behavior
  across a period?
• Bonding between atoms changes across a period
  
• Bonding involves the valence electrons or
  outermost shell (or highest shell) electrons
• Atoms form bonds to become more stable electrons
  are gained, lost or shared to achieve stability.
• The properties of a compound are different from
  the properties of the atoms that make up the
  compound. Ex NaCl

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Types of Bonds

• 1. Ionic bond
• Transfer of e- from a metal to a nonmetal and
  the resulting electrostatic force that holds them
  together forms an ionic compound.
• EX Na Cl- ? NaCl
• (neutral)

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Ionic Bonding

• Ionic bonds involve the formation of positive and
  negative ions that then attract each other.
• Metals form positive ions by losing electrons
• Nonmetals form negative ions by gaining electrons

Next Slide
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Ionic Bonding Example 1

• Sodium has 1 valence electron which it needs to
  lose.

Chlorine has 7 valence electrons and needs to
pick up 1 electron.
Next Slide
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Ionic Bonding Example 1

• The sodium loses its electron to the chlorine.

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-1
This makes the sodium 1 and the chlorine -1
They attract each other forming the compound
NaCl
Next Slide
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Ionic Bonding Example 2

• Magnesium has 2 valence electrons which it needs
  to lose.

Oxygen has 6 valence electrons, It needs to pick
up 2 electrons.
Next Slide
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Ionic Bonding Example 2

• Magnesium loses both of its outer electrons to
  the oxygen.

Next Slide
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Ionic Bonding Example 2

• This gives the magnesium a 2 charge

and the oxygen a -2 charge
They join together to form the compound MgO.
Next Slide
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Exchange of Electrons
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Ionic Bonding

• When atoms bond, the properties of the new
  compound are DIFFERENT from the properties of the
  elements that made them up.
• Ionic compounds have several characteristics in
  common due to the presence of the ionic bond.
• These characteristics include
• Crystalline structure (the formula gives the
  ratio between the ions making up the substance)
• High melting points, making them solids at room
  temperature
• Usually water soluble (can dissolve in water)
• Electrolytes when in solution (conduct
  electricity)

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Ionic Bonding

• Sodium chloride (NaCl) is held together by an
  ionic bond.
• The properties of sodium chloride are
• Sodium chloride forms a cube shaped
  crystalline solid.

Melting point 801C Boiling point
1413C Highly soluble in water Strong electrolyte

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Types of Bonds

• 2. Covalent bond
• Formed from the sharing of e- pairs between two
  or more nonmetals resulting in a molecule.
• EX H2 O ? H2O

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Covalent Bonding

• Definition - bond formed due to the sharing of
  electrons between nonmetals.
• The high attraction for electrons of nonmetals
  results in the nonmetals attempting to remove
  electrons from each other.
• Since neither nonmetal is able to give up
  electrons they are forced to share the electrons.

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Covalent Bonding Example 1

• Bromine and Fluorine both have 7 valence
  electrons and very high attraction for electrons.

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Covalent Bonding Example 1

• Since neither fluorine or bromine are able to
  lose electrons they get drawn together until
  their outer orbits overlap and one electron from
  each atom goes back and forth between the two
  atoms.

This creates the compound BrF.
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Covalent Bonding Example 2

• Hydrogen and Oxygen can both pick up electrons.
  (If hydrogen loses its only electron it will end
  up as a nucleus with no electrons around it.)
  They will share electrons to form covalent bonds.
• This results in the formula H2O

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Covalent Bonding Example 3

• Seven of the elements have such high attraction
  for electrons that they will never exist as
  individual, unattached atoms. Anytime these
  elements are present in pure form they will bond
  to other atoms of the same element.
• For example a fluorine atom will readily bond to
  a second fluorine atom.
• Resulting in F2.

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Covalent Bonding

• These elements are called diatomic elements, and
  the molecules they form are called diatomic
  molecules. The definition of a diatomic molecule
  is
• A molecule made up of 2 atoms of the same element.

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Covalent Bonding

• The seven diatomic elements are
• Hydrogen, H2
• Nitrogen, N2
• Oxygen, O2
• Fluorine, F2
• Chlorine, Cl2
• Bromine, Br2
• Iodine, I2

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Covalent Bonding

• Covalent compounds are made up of small units
  called molecules. The formula for a covalent
  compound tells the actual number of atoms, of
  each element, found in each molecule.
• For example The formula for water is H2O. This
  formula indicates that each water molecule is
  made up of two Hydrogen atoms and one Oxygen atom.

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What is a polar covalent bond?

• Covalent bonds involve sharing electrons.
• The electrons may be shard equally (nonpolar
  covalent) or unequally (polar covalent).
• Example H2 shares electrons equally, but HCl
  does not. Therefore, H2 contains a nonpolar
  covalent bond and HCl contains a polar covalent
  bond.

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Polarity of water

• Take for example H2O. When we draw the structure
  it looks like
• O
• H H
• The oxygen atom pulls electrons away from the
  hydrogen atoms. This unequal sharing results in
  polar bonds, which have a more negative end and
  a more positive end.

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Polar Water Molecule
more negative

• EX O
• H H
• H2O is a polar molecule.
• Bond polarity affects the properties of a
  material such as melting and boiling points,
  crystal structure and acidity.

more positive
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Covalent Bonding

• The characteristics shared by covalent compounds
  are
• Molecular structure individual units
• Low melting and boiling points, most covalent
  compounds are gases or liquids at room
  temperature (the larger the molecule the higher
  its melting and boiling point)
• Soluble in covalent solvents such as alcohol or
  benzene.
• Nonelectrolytes

gas atoms and molecules
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Covalent Bonding

• Paradichlorobenzene (moth balls) (C6H4Cl2) is a
  covalent compound.
• Its properties are
• Molecular solid

Melting point 53.1C Boiling point
174.55C Soluble in alcohol, ether, acetone
and benzene Nonelectrolyte
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Comparison of Bonding Types
ionic
covalent
ions
molecules
The properties of a material depend on the
structure -different bond types result in
different properties.
molten salts conductive
non- conductive
Both determined by valence electrons
transfer of electrons
sharing of electrons
high mp
low mp
not usually water soluble
water soluble
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Polyatomic Ions/Radicals

• Some groups of atoms are covalently bonded
  together so strongly that the stay together
  during chemical reactions and act as a single
  unit.
• These groups of atoms become charged, with the
  charge being spread out through out the group.
• These groups are called polyatomic ions or
  radicals.

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Polyatomic Ions/Radicals

• Definition - groups of atoms bonded together that
  act as a charged unit.
• Examples
• Ammonium - NH41
• Sulfate - SO4-2
• Acetate - C2H3O2-1
• Phosphate - PO4-3
• What kind of bond do you think hold the atoms in
  a polyatomic ion together?
• What kind would hold two polyatomic ions together?

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Type of bond? Ionic, Polar Covalent, or
Nonpolar Covalent?
TiO2
CH4
NaI
CS2
O2
KCl
AlCl3
CsF
HBr
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Types of Bonds

• 3. Metallic bond
• Metals bonding with other metals do not gain or
  lose e- or share e- unequally. These bonds are
  created from the delocalized e- that hold
  metallic atoms together.

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Chemical Formulas

• A chemical formula is a combination of symbols
  that represents the composition of a compound.
• Chemical symbols are used to indicate types of
  elements present.
• Subscripts are used to indicate the number of
  atoms for each element present.

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What are the parts of a formula?

• chemical symbols
• C8H18
• number of atoms of each element
• 8 atoms of carbon
• 18 atoms of hydrogen

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Charges of Monatomic Ions

• Because atoms want to reach an octet of valence
  electrons, the oxidation numbers, (positive or
  negative charges) can be predicted for single
  atoms (monatomic).
• Metals tend to have positive oxidation numbers.
  (lose e-)
• Nonmetals tend to have negative oxidations
  numbers. (gain e-)

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Remember...
0
1
2
3
varies
3-
2-
1-
varies 1 to 7
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Oxidation Numbers (charges) of Polyatomic Ions

• Polyatomic ions are ions that are made up of two
  or more atoms.
• Refer to your table of Polyatomic ions.
• Polyatomic ions generally have the following
  endings ate or ite
• Ex
• NO2- nitrite
• PO43- phosphate

SO42- sulfate
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Polyatomic Ions
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Oxidation Numbers

• The sum of the oxidation numbers in a compound
  must equal zero.
• Ex CaCl2 Ca2 Cl- Cl-
• 2 positive charges
• 2 negative charges 0
• The charge on a monatomic ion is its oxidation
  number.
• Ex Ba2 has an oxidation of 2
• Cl- has an oxidation of -1

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What happens when the predicted charge can vary??

• The oxidation number of a transition element is
  shown using Roman numerals to indicate the
  charge.
• The Roman numeral indicating oxidation.
• Ex iron (II) is Fe2
• iron (III) is Fe3

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Why is aluminum oxide Al2O3?
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Writing Ionic Formulas

• Ionic compounds are composed of metals and
  nonmetals.
• Ionic compounds are made from the gaining or
  losing of electrons and the resulting
  electrostatic force that holds the ions together.
  
• The sum of the oxidation numbers in a compound
  must equal zero.

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Writing Ionic Formulas

• When writing formulas, the cation (metal ion) is
  always written before the anion (nonmetal ion).
• When using polyatomic ions, refer to charge given
  on your table.
• NOTE There is only one polyatomic
• cation (NH4).
• The rest are all are polyatomic anions.

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Criss Cross Method of Writing Formulas

• Notice a trend between the oxidation/charges of
  ions and the subscripts of elements.
• Ex Mg2 and Cl- gives
• MgCl2
• Ex Al3 and SO4-2 gives
• Al2(SO4)3

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Criss Cross Method of Writing Formulas

• This method crosses charges and subscripts to
  form neutral compounds.
• Al 3 and O-2
• Al and O
• Al2O3
• (neutral)

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Criss-Cross Method of Writing Formulas

• Ex
• Lead (II) phosphate
• Pb2 and PO4-3
• Pb and PO4
• Pb3(PO4)2

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Nomenclature

• Nomenclature is defined as a naming system.
• Chemistry uses nomenclature to standardize names
  of chemicals.
• Lets take a look.

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Naming Binary Ionic Compounds

• Rules for naming binary (composed of two) ionic
  compounds
• 1. Name of cation is given first. The name of
  the cation is the same as the element.
• 2. Name of anion is given last. The name of the
  anion is the same as the element, but with an
  ide suffix.

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Naming Ionic Compounds

• Ex Al2O3
• Aluminum and Oxygen
• Aluminum oxide
• cation anion

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Naming Ionic Compounds

• Ex Ni2O3
• Note that the cation has MORE THAN ONE possible
  oxidation state, so Roman numerals are needed to
  identify the ion.
• Nickel and Oxygen
• Nickel (III) oxide

-2
3
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Naming Ionic Compounds

• EX AgCl is __________________
• Silver chloride
• Na2O is __________________
• Sodium oxide
• CaBr2 is ___________________
• Calcium bromide
• PbO2 is ___________________
• Lead (IV) oxide

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Naming Ionic Compounds with a polyatomic ion

• Rules for naming compounds that contain a
  polyatomic ion.
• 1. Cation rule from binary applies.
• 2. Anion takes the name of the polyatomic
  ion as found on the table.
• Ex Al2(SO4)3
• aluminum sulfate
• Ex Mg(OH)2
• magnesium hydroxide

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Naming Ionic Compounds

• Ex Li2CO3 is ____________________
• Lithium carbonate
• Ba(OH)2 is ___________________
• Barium hydroxide
• Zn(NO3)2 is __________________
• Zinc nitrate
• KClO3 is ___________________
• Potassium chlorate

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Naming continued

• Name the following compound
• Ba(Na)2
• Banana

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Naming Binary Molecular Compounds

• Unlike ionic compounds, molecular compounds are
  composed of individual covalently bonded units,
  or molecules.
• Covalent compounds are formed between nonmetals.
• Prefixes are used to indicate number of each type
  of element in the compound.
• Write the prefixes as indicated on the next
  slide.

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Molecular Prefixes
1 mono- 6 hexa-
2 di- 7 hepta-
3 tri- 8 octa-
4 tetra- 9 nona-
5 penta- 10 deca-
(decade)
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Naming Binary Molecular Compounds

• Follow these rules
• The element written first is given a prefix if it
  contributes more than one atom to the molecule.
• The second element is named by combining (a) a
  prefix indicating the number of atoms contributed
  by the element, (b) the root of the name of the
  second element, and (c) the ending -ide.
• 3. The o or a at the end of a prefix is
  usually dropped when the word following the
  prefix begins with another vowel.
• Ex monoxide or pentoxide

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Naming Binary Molecular Compounds

• Ex P4O10
• P has more than one atom in this molecule.
• Tetraphosphorus
• O is named by combining prefix, root name, and
  -ide ending.
• decoxide (a is dropped from prefix)
• combine to form Tetraphosphorus decoxide

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Naming Binary Molecular Compounds

• Ex SO3 is ____________________
• sulfur trioxide
• PBr5 is ____________________
• phosphorus pentabromide
• ICl3 is _____________________
• iodine trichloride
• H2O is _____________________
• dihydrogen monoxide
• Sb2O3 is _____________________
• diantimony trioxide (metalloid)

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• THE END

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Strange Names for Molecules

• BUCKMINSTER MORONIC ACID FULLERENE

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Chromatography Lab
Paper Chromatography Paper chromatography is a
method chemists use to separate compounds from
one another, but not change them. In this lab we
will explore how this separation is made using
different dye compounds. Molecules with similar
polarities or molecular structures are attracted
to each other. Water molecules have a polar
structure. Because of this structure the oxygen
end of the molecule has a small negative
electrical charge and the hydrogen end has a
small positive charge. Liquid water is held
together by the attraction between the charges on
different molecules.
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• A more complex, yet still similar molecule is
  cellulose, a molecule which is the basic
  component of paper. It is a very long molecule (a
  polymer) in which thousands of rings of six atoms
  each are linked together like beads. A portion of
  a cellulose molecule is shown below.

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• Paper Chromatography Paper chromatography is a
  method chemists use to separate compounds from
  one another, but not change them. The polar
  regions of these molecules are attracted to polar
  regions of the cellulose chains (which help to
  hold the fibers together in paper). Not
  surprisingly, water molecules, being polar, are
  also attracted to these regions and when paper is
  wet it loses strength because the water molecules
  get between the cellulose chains and weaken the
  attraction between them.
• When water molecules move up paper that is dipped
  in water, the molecules which might be dissolved
  in the water will also be carried along up the
  paper. This is applied to the separation of dyes
  in a technique known as paper chromatography.

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• A spot of dye is placed on the paper above the
  level of the water. As the water moves up, the
  dye molecules will move with it if they are more
  strongly attracted to the water molecules than to
  the paper molecules. If the dye molecules are
  more strongly attracted to the paper than to the
  water, they will move more slowly than the water
  or even not at all.
• What if the dye is a mixture? If two or more dyes
  have been mixed, then they may move at different
  rates as the water moves up the paper. If this
  happens, they will separate and we can identify
  them . This is depicted in the sketches below.

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Yellow 5 Structure

• C16 H9 N4 Na3 O9 S2

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• Skittles ingredients sugar, corn syrup,
  hydrogenated palm kernel oil, apple juice from
  concentrate, citric acid, dextrin, natural and
  artificial flavors, gelatin, food starch,
  coloring (includes Yellow 6 lake, Red 40 lake,
  Yellow 5 lake, Blue 2 lake, Blue 1 lake, Yellow
  5, Red 40, Yellow 6, Blue 1), ascorbic acid
  (Vitamin C).
• MM ingredients Milk Chocolate (Sugar,
  Chocolate, Cocoa Butter, Skim Milk, Milkfat,
  Lactose, Soy Lecithin, Salt, Artificial Flavors),
  Sugar, Cornstarch, Corn Syrup, Dextrin, Coloring
  (Includes Blue 1 Lake, Red 40 Lake, Yellow 6,
  Yellow 5, Red 40, Blue 1, Blue 2 Lake, Yellow 6
  Lake, Yellow 5 Lake, Blue 2), Gum Acacia.

The lake part means that the dye is attached
as a coating