Transcript CH. 6

CH. 6
The Structure of Matter
Ch. 6 Section 1 Notes
• Compounds and Molecules
– Pg. 177-182
Chemical Bonds
• The forces that hold atoms or ions
together in a compound are called
chemical bonds.
– Can be broken, and the atoms rearrange
Chemical Structure
• The structure of a building is the way the
building’s parts fit together
– A compound’s chemical structure is the way
the atoms are bonded to make the compound
• Some models represent bond lengths and
angles.
– Bond length is the distance between the
nuclei of two bonded atoms
– If a compound has 3 or more atoms, a bond
angle (the angle formed by two bonds to the
same atom) tells which way there atoms
point.
– Atoms are often represented by a ball-and-stick model to help
you understand the compounds structure.
– Structural formulas also show the structures of compounds.
• Chemical symbols are used to represent the atoms
– Space-filling model is another way to represent a water
molecule.
• Shows the space that the oxygen and hydrogen atoms take
up, or fill
• The chemical structure of a compound
determines that properties of that
compound.
– Compounds with network structures are
strong solids.
• Ex: quartz
– The strong bonds make the melting and boiling point of
quartz and other minerals very high
• Some networks are made of bonded ions.
– Ex: Table Salt (NaCl)
• Found in the form of regularly shaped crystals
• Made of a repeating network connected by strong
bonds
• Oppositely attracted ions
• High melting and boiling point
• Some compounds are made of molecules.
– Ex: sugar
• Molecules attract each other and form crystals
Nitrogen, Oxygen, and Carbon Dioxide
--Gases that are made of molecules
--Atoms are strongly attracted to each other and
are bonded
• The strength of attractions between
molecules varies.
• Sugar, water and Dihydrogen sulfide are all
compounds made of molecules but have different
properties
*The higher the melting point, the stronger the
attraction between the atoms
• Hydrogen Bond
– Oxygen atom of a water molecule is attracted
to a hydrogen atom of another molecule
• Strong bonds within each water molecule
• Weaker attractions between water molecules
Ch. 6 Section 2 Notes
• Ionic and Covalent Bonding
– Pg. 183-190
Why do Chemical Bonds Form?
• Atoms join to form bonds so that each
atom has a stable electron configuration.
– One similar to a noble gas
• There are two kinds of chemical bonding:
– Ionic Bonding
– Covalent Bonding
Ionic
Compounds
Covalent
Compounds
Structure
Network of
bonded ions
Molecules
Valence
Electrons
Transferred
Shared
Electrical
conductivity
Good (when
melted or
dissolved)
Poor
Solid
Solid, liquid, or
gas
Generally high
Generally low
State at room
temp.
Melting and
boiling points
Ionic Bonds
• Form from the attractions between such oppositely
charged ions.
• Formed by the transfer of electrons
– Oppositely charged ions bond (NaCl)
• Ionic compounds are in the form of networks, not
molecules.
• The formula unit of one sodium ion and one chloride ion
is NaCl
– NaCl ratio is 1:1
– CaF2 is 1:2
• When melted or dissolved in water,
ionic compounds conduct electricity.
– Ions are free to move
when not is solid form.
Covalent Bonds
• Compounds that are made of molecules,
such as water and sugar have covalent
bonds.
• Atoms joined by covalent bonds share
electrons.
• Usually form between nonmetal atoms.
• Can be solid, liquid or gas
• Low melting points
• MOST do not conduct electricity (not
charged)
• Example: Cl2
– Each has 7 valence electrons.
• Share one electron to have 8 valence electrons
and become stable.
Atoms may share more than one pair of
electrons.
• When drawing the electron dot diagram, a line — means
that there are 2 electrons being shared.
• If there is two lines ==, that is a double covalent bond (4
electrons being shared)
• A triple covalent bond is
formed by bonding two
nitrogen atoms
(total of 6 electrons)
Atoms do not always share
electrons equally.
• When electrons are shared equally, they are called
nonpolar covalent bonds.
– Ex: Cl2
• When two atoms of different elements share electrons,
the electrons are not shared equally and forms a polar
covalent bond.
– Ex: NH3
Metallic Bonds
• Metals are flexible and conduct electric
current well because their atoms and
electrons can move freely throughout a
metal’s packed structure.
• Atoms in metals such as copper form
metallic bonds.
Polyatomic Ions
• Acts as a single unit in a
compound, just as ions that
consist of a single atom do.
• Groups of covalently bonded
atoms that have a positive or
negative charge as a group.
– Both covalent and ionic bonds
• There are many common
polyatomic ions.
Parentheses group the atoms of a
polyatomic ion.
• A polyatomic ions charge
applies not only to the last atom
in the formula but to the whole
ion.
• A polyatomic ion acts as a
single unit in a compound
• Some names of polyatomic anions relate
to the oxygen content of the anion.
• Most end with –ite or –ate
– -ate ending usually used to name an ion that
has 3 oxygen atoms
• Examples: sulfate (SO42–), nitrate (NO3–),
chlorate (ClO3–)
– 2 or less oxygen atoms have an –ite ending
• Examples: sulfite (SO32–), nitrite (NO2–),
chlorite (ClO2–)
• Hydroxide and Cyanide are exceptions to
the rules.
CH. 6 Section 3 Notes
• Compound Names and Formulas
– Pg. 191-196
Naming Ionic Compounds
• Formed between cations and anions
• The names of ionic compounds consist of
the names of the ions that make up the
compounds.
• Names of cations include the elements of
which they are composed.
– Usually the name of the element
• Ex: sodium forms a sodium ion
• Names of anions are altered names of
elements.
– The difference is the name’s ending
• Usually with the ending –ide
• Compounds with Oxygen atoms have –ate, or –ite
endings
• An ionic compound must have a total
charge of zero.
• Some cation names must show their
charge.
– Transition metals may form several cations
(each will have a different charge).
– Iron forms a +2 ion AND a +3 ion
• This is shown by placing the charge of the cation
as a Roman numeral in parentheses.
– Iron (II) ion and Iron (III) ion
– FeO --- Iron (II) Oxide
– Fe2O3 --- Iron (III) Oxide
Determining the charge of a
transition metal cation.
• The total charge of the compound MUST
be zero.
– Fe2O3
• Three oxide ions have a total charge of 6-. (each
oxygen ion has a charge of 2- 2-(3)=6-)
– So, the total charge of the cation must be 6+
Writing Formulas for Ionic
Compounds
• If you are given the compound’s name:
you can find the formula
• If you are given the formula: you can find
the charge of each ion
Naming Ionic Compounds Rules
• If you are given the Name:
– 1. Find the symbol of each
element
– 2. Find the charge of each ion
– 3. Criss-cross Method
– 4. If one of the ions is a
Polyatomic Ion, put
parentheses around it!!
• Calcium Chloride
1. Ca, Cl
2. Ca+2 , Cl -1
3. CaCl2
4. This is not a
polyatomic
Ion
5. This is a
polyatomic
Ion
Naming Ionic Compound Rules:
• If you are given the formula:
–
• AgF
1. Since there is no
1. Determine if the FIRST ion
subscript number
the charges for
is a Transition metal. If so, you
both must be 1.
MUST find it’s charge!
– 2. Find the name of each of the
ions
– 3. The cation is the same as it
is on the periodic table
– 4. The anion has an –ide
ending (unless it is a
polyatomic ion)
2. Ag is Silver, F is
Flourine.
3. F is in group 17
and has a -1
charge so, Ag is
the cation.
4. Silver Flouride
• To find the charge of ions in a
chemical formula:
1. Determine the ratio of the given
formula
2. Separate the ions
3. Determine each of their charges
4. If the cation is a transition metal, use
the criss-cross method and then look
at it’s ratio.
5. Compare to the original ratio. What
ever you do to the first element, you
must do the the 2nd.
CrO2
1. 1:2 ratio
2. Cr O2
3. Cr +4 O-2
4. Cr2O4 Ratio is
2:4
5. Reduce the
ratio to 1:2
Math Skills “Writing Ionic Formulas”
Practice Problems 1-3 Pg. 193
1. Lithium oxide
1. Li+1 O-2
2. Li2O
2. Beryllium chloride
1. Be+2 Cl-1
2. BeCl2
3. Titanium (III) nitride
1. Ti+3 N-3
2. TiN
Naming Covalent Compounds
• For covalent compounds of two elements, numerical
prefixes tell how many atoms of each element are in the
molecule.
• Numerical prefixes are used to name covalent
compounds of two elements.
– If there is only one atom of the first element, the name
does not get a prefix.
Number of
Atoms
1
2
3
4
5
6
7
8
9
10
Prefix
MonoDiTriTetraPentaHexaHeptaOctaNonaDeca-
• BF3
– Boron Trifluoride
• N2O4
– Dinitrogen tetroxide
Empirical Formulas
• Chemical formulas that are unknown are determined by
figuring out the mass of each element in the compound.
• Once the mass of each element is known, scientists can
calculate the compound’s empirical formula, or
simplest formula.
• An empirical formula tells us the smallest whole-number
ratio of atoms that are in a compound.
• Different compounds can have the same
empirical formula.
• Molecular formulas are determined from
empirical formulas.
• A compound’s molecular formula tells you how
many atoms are in one molecule of the
compound.
• Masses can be used to determine the empirical
formula.
– Convert the masses to moles. Then, find the molar
ratio to give you the empirical formula.
• Pg.196
– Math Skills “Finding Empirical Formulas”
1. One mole of an unknown compound has 36.04 g of
Carbon and 6.04g of hydrogen. What is the
compound empirical formula.
Section 3 Review # 1, 5
1. Name the following ionic
compounds, and specify
the charge of any
transition metal cations.
a)
FeI2
a) Iron(II) Fluoride
b) MnF3
a) Manganese(III)Flouride
c) CrCl2
a) Chormium(II) Chloride
d) CuS
a) Copper(II) Sulfide
5. Determine the
chemical formulas for
the following ionic
compounds.
a)
Magnesium sulfate
a) MgSO4
b) Rubidium bromide
a) RbBr
c) Chromium(II) fluoride
a) CrF2
d) Nickel(I) carbonate
a) Ni2CO3
Ch. 6 Section 4 Notes
• Organic and Biochemical Compounds
– Pg. 197-204
Organic Compounds
• An organic compound is a covalently
bonded compound that contains carbon.
– Most contain hydrogen.
– Oxygen, nitrogen, sulfur, and phosphorus can
also be found in organic compounds.
• Carbon atoms form four covalent bonds in
organic compounds.
• A compound made of only hydrogen and
carbon atoms is known as a hydrocarbon.
– Methane, CH4 is an example
• There are four single C-H bonds
• A carbon atom may never form more than 4 bonds
at a time.
• Alkanes are hydrocarbons that have only
single covalent bonds.
– Can have C-C bonds as well as C-H bonds
– Methane is the simplest alkane
Arrangements of carbon atoms in
alkanes.
• The carbon atoms in methane, ethane,
and propane are all bonded in a single line
because that is their only possible
arrangement.
• If there are more than 3 bonded carbon
atoms in a molecule, the carbon atoms do
not have to be in a single line.
• IF they are in a single line: the alkane is a
normal alkane, or n-alkane.
• The condensed structural formula shows
how the atoms bond.
• Alkane chemical formulas usually follow a
pattern.
– Except for cyclic alkanes
• The # of Hydrogen atoms is always 2
more than 2x the # of carbon atoms
– CnH2n+2
Alkenes have double carboncarbon bonds.
• Hydrocarbons
• Have at least one double covalent bond
between carbon atoms. C=C
• Replace the –ane ending with –ene.
• Simplest alkene is ethene (ethylene)
• Alcohols have hydroxyl (-OH) groups.
• Made of oxygen, hydrogen, and carbon
• Most alcohols end in –ol
• Alcohol and water molecules behave similarly.
• Methanol and methane are alike except that one
of the hydrogen atoms is replace by a Hydroxyl
group
• Alcohol molecules are attracted to each other
– Liquid at room temp; HIGH boiling points
Polymers
• A polymer is a molecule that is a long
chain made of smaller molecules.
– Have repeating subunits
– Polyethene, is a polymer that makes up
plastic milk jugs.
“Poly”=many
• Ethene is an alkene that has the formula C2H4.
– Polyethene means “many ethenes”
– The smaller molecule that makes up the polymer is
called a monomer.
• Some polymers are natural, and others
are artificial.
• Natural: Rubber, wood, cotton, wool,
starch, protein, DNA, etc.
• Human-made: Plastics or Fibers
A polymer’s structure determines
its elasticity.
• Chains are tangled and can slide past
each other.
• When the chains are connected to each
other, the polymer’s properties are
different.
– Some are elastic (can stretch)
• When released, returns back to its original shape.
– Ex: Rubber bands
Biochemical Compounds
• Essential to life, include carbohydrates, proteins,
and DNA
• Can be made by living things
• Carbohydrates give you energy
• Proteins form important parts of your body
– Muscles, tendons, fingernails, and hair
• The DNA inside your cells gives your body info
about what proteins you need.
Many carbohydrates are made of
glucose.
• Carbohydrates include sugars and starches,
provide energy to living things.
• Sucrose (table sugar) is made of two simple
carbohydrates, glucose and fructose, bonded
together.
• Starch is made of a series of bonded glucose
molecules, and is a polymer.
• When you eat starchy food the enzymes in
your body break down the starch.
• The glucose that is not needed is stored
as glycogen, a polymer of glucose.
• When active, glycogen breaks apart into
glucose molecules and gives you energy.
Proteins are complex polymers of
amino acids.
• Proteins, which provide structure and function to parts
of cells, are very complex.
• Made of many different molecules that are called amino
acids.
– Made of carbon, hydrogen, oxygen, and nitrogen.
Some contain sulfur.
– 20 amino acids found in naturally occurring proteins
• The amino acids that make up a protein
determine the protein’s structure and
function.
• Proteins are long chains made of amino
acids.
– Made of thousands of bonded amino acid
molecules
DNA is a polymer that stores
genetic information
• DNA is a very long molecule made of carbon,
hydrogen, oxygen, nitrogen, and phosphorus.
• DNA is in the form of paired chains, or strands.
– Shape of a twisted ladder, double helix.
• DNA is the information that the cell uses to
make proteins.
• DNA monomers:
– Adenine, thymine, cytosine, and guanine
• Pair with other DNA monomers that are
attached to the opposite strand in a
predictable way