Transcript Chemistry Notes for the Whole Year Powerpoint presentation

Atomic Structure
• What is an atom?
• An atom is the smallest unit of matter. It contains
protons (+ charge), neutrons (0 charge), and
electrons (- charge). The protons and neutrons are
located in the nucleus of the atom. Electrons orbit
around the nucleus in distinct energy levels.
• The nucleus of the atom is small, dense, positively
charged, and heavy. Ernest Rutherford determined
this during his gold foil experiment.
Rutherford Experiment
Rutherford Experiment
Atomic Structure
• The nucleus is very small compared to the rest of
the atom. Its size is comparable to a marble
resting at the 50 yard line of Qualcomm Stadium.
Therefore, the atom is mostly made up of empty
space.
• Even though like charges repel, the protons in the
nucleus do not repel because they interact with
and are spaced out by the neutrons.
Protons, Neutrons, and Electrons
• How do we find the numbers of each
subatomic particle?
• The number of protons=atomic number (the
whole number on the periodic table). Do
not confuse atomic number with atomic
mass, they are two different things.
• #protons=#electrons in an atom with no
charge.
# neutrons
• Mass number= #protons + #neutrons. This can be
rearranged to determine neutrons. How can we do
it?
• The mass number cannot be found on the periodic
table. The atomic mass is the average of an
element’s mass numbers, but not the mass number
we use for problems.
• How many protons, neutrons, and electrons are in
phosphorous 31? (mass number=31)
Atom Example
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Problems
• How many protons, neutrons and electrons
are in Oxygen 16 (mass number =16)?
• How many protons, neutrons, and electrons
are in Iron 56 (mass number=56)?
What are isotopes?
• Isotopes are atoms of the same element that have
different numbers of neutrons.
• Some isotopes are stable and some are not. The
unstable isotopes decay and we call them
radioactive. Unstable isotopes decay at a known
rate and allow us to date artifacts and rocks using
that known rate.
• Consider Carbon 12, Carbon 13, and Carbon 14.
Are these isotopes? How many protons, neutrons,
and electrons are found in each?
• I will show you isotope notation on the board also.
Isotope Examples
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What are periodic families?
• The rows on the periodic table are called
periods. Each period represents an energy
level. The columns are called groups.
• The families on the periodic table occur in
columns (up and down) and contain
elements that have the same physical and
chemical properties.
• Elements that are to the right of the staircase are nonmetals. Elements that are to the left of the staircase are
metals.
• Elements in column one are called “alkali metals.”
• Elements in column two are called “alkaline earth metals.”
• Elements in columns 3-12 are called “transition metals.”
• Elements in column 17 are called “halogens”
• Elements in column 18 are called “noble gases” or “inert
gases.”
• Silicon and Germaniuim are “metalloids” meaning that
they have properties of both metals and non-metals.
• What families do these
elements belong in?
Questions
• Name an element that is a member of:
The halogens
The inert gases
The transition metals
The alkali metals
• Mg is a member of what family?
What are the periodic trends?
• When we look at periodic trends, across means we
go from left to right and down a group means that
we go from top to bottom. These trends are not
readily obvious from the periodic table.
• Atomic size/ionic size- This is the size of the
atom or ion. This decreases across a row and
increases when you go down a group.
• Electronegativity-this is the tendency of an
element to attract electrons. This increases across
a row and decreases down a group.
• Ionization energy-this is the amount of energy
required for an atom to eject one of its electrons.
This increases across a row and decreases down a
group.
• Between Al, P, and Cl, which element has the
higher electronegativity?
Questions
• Between N, P, and As, which element has
the highest atomic size?
• Between C, N, and O, which element has
the highest ionization energy?
What is radioactive decay?
• Radioactive decay occurs when an unstable
isotope decays into smaller isotopes.
• There are three types of radioactive decay. Alpha
()-this is when a helium nucleus is ejected from
an isotope. This is the weakest form of radiation.
Paper and clothing can stop it.
• Beta()-this occurs when a neutron emits an
electron. This neutron becomes a proton. This is
middle strength radiation and wood or cardboard
can block it.
Penetrating Power of Radiation
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Types of radiation
• Gamma()- this is high energy radiation that
a radioactive isotope can emit when it
decays. It causes no change in the nucleus.
However, it is the strongest form of
radiation. Lead and concrete can block it.
• For nuclear reactions, please look on the
board.
What is fission or fusion?
• Fission is when a radioactive isotope splits apart into smaller
components. A lot of energy and gamma radiation is released.
• Fusion is when two hydrogen nuclei fuse together to create a helium
nucleus. A lot of energy and gamma radiation is released (more than
fission).
• Fusion of higher elements occurred in dying stars. There was enough
heat and pressure, within those dying stars, for the helium to fuse with
other helium atoms and form larger elements. The destruction of that
star then diffused those elements into the universe.
• There was also enough energy for those elements, once they had
formed, to fuse with other nuclei and form larger elements in dying
stars.
• Thus, all life on this planet is the product of dying stars.
Examples of Fusion and Fission
• Fission is first picture.
• Fusion is in second
picture.
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Quic kT i me™ and a
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What are orbitals?
• Orbitals are clouds where electrons reside in an atom.
• Each orbital can hold a maximum of two electrons (one
with up spin and one with down spin).
• The types of orbitals are:
s (1 orbital, energy level=row # on periodic table)
p (3 orbitals, energy level=row # on periodic table)
d (5 orbitals, energy level=row # on periodic table -1 )
f (7 orbitals, energy level=row # on periodic table -2 )
Orbital example
Question?
• How many electrons can each orbital hold?
• Where are these orbital blocks located on
the periodic table?
What are electron configurations?
• Electron configurations show how electrons are arranged around an
atom.
• Three rules govern electron configurations and they are:
• Aufbau principle-electrons enter orbitals of lowest energy first. For all
electron configurations, start filling electrons in the 1s orbital, which is
the lowest energy level.
• Pauli exclusion principle-An atomic orbital may describe at most two
electrons (no more than two electrons per orbital)
• Hund’s rule-When electrons occupy orbitals of equal energy, one
electron enters each orbital until all the orbitals contain one electron
with spins parallel. Second electrons then add to each orbital so that
their spins are paired with those of the first electrons.
How do we write electron
configurations?
• Electron configurations contain the energy level, orbital, and number
of electrons in that orbital. For example, 1s1, the first number means
energy level one, s is the orbital, and the superscript number means
that there is one electron in this orbital.
• Start at hydrogen and fill electrons from left to right starting from
there. Count how many electrons are in these orbitals as you move
from left to right. Also, make sure you change to the correct orbital
once you move to another orbital’s block on the periodic table.
• You can move to another row on the periodic table after you have
moved to all the way over to the right on the previous row. Remember
that energy levels change when you move to another row on the
periodic table. Do this until you have all electrons accounted for in the
element.
Write electron configurations
• Write the electron configuration for hydrogen.
• Hydrogen has one electron. Begin at row one, s
block, and count over one. That is hydrogen.
• From that, we are in energy level one, the s orbital,
and there is only one electron in it. So we write,
1s1 for the electron configuration.
• What about for helium? Write its electron
configuration.
Orbital Blocks
More electron configurations
• Write the electron configuration for Li.
• It should be:
• 1s2 2s1 and this shows three total electrons, which
lithium has.
• Write the electron configuration for N
• It should be:
• 1s2 2s2 2p3
Write the electron configuration
for Fe
How do we write short-hand
electron configurations?
• Start short-hand electron configurations with the nearest
noble gas that has fewer electrons than the element you are
looking at.
• Write that noble gas and its number of electrons once you
begin your electron configuration. Then, write your
electron configuration from that noble gas.
• For example, for Mg, you would begin its short-hand
electron configuration with [Ne]10
• That would take care of the first 10 electrons in that
configuration (through row 2). You would then move to
row three and complete the electron configuration.
• It would look like: [Ne]10 3s2
Write the short-hand electron
configuration for:
•
•
•
•
Na
Fe
Br
As
What is another way to write
electron configurations?
• This method of writing electron configurations shows the orbital (1s
for example) and how many electrons are in it. Arrows represent the
electrons (one arrow points up meaning that the electron has up spin,
and another arrow points down meaning that the other electron has
down spin) and are placed in circles that represent the orbitals.
• When writing electron configurations in this way, follow Hund’s rule
meaning that you must fill each orbital, in a particular subshell, with
one electron that has the same spin (usually an up arrow). After filling
these orbitals, then you would place the second electron in each orbital
and it would have down spin (opposite of the first electron).
• For example, for the p orbital, you would place one electron (up arrow)
in each of the three orbitals before filling each orbital with the second
electron. The same holds true for the d orbitals.
Example
• Look on the whiteboard for examples of this
style of electron configuration.
• Write electron configurations for the
following:
Ne
O
Na
How do we use the short-hand
method?
• For this type of electron configuration, you can
use the short-hand method. You find the nearest
inert gas (to the element) that has fewer electrons
and place it in brackets at the beginning of the
configuration. Then you write the configuration
from there showing the orbitals and how many
electrons are in them.
• Write short-hand electron configurations for:
Fe
Br
Sr
What are valence electrons?
• Valence electrons are electrons in the outermost energy
level of an atom. They are the electrons that are involved
in bonding of atoms.
• Core electrons are electrons that are in the inner energy
levels of an atom. They are not involved in bonding.
• s and p electrons (if they are in the same highest energy
level) automatically count as valence electrons.
• d electrons count as valence electrons if there are 5 or
fewer of them in the outermost energy level. If there are 6
or more d electrons then they do not count as valence
electrons.
How do we determine the
number of valence electrons?
• We must first find the highest energy level in an atom
(look at the highest energy number in its electron
configuration, this is easy to do if you use short-hand).
Count the number of electrons in that highest energy level
(s and/or p electrons if they are in the same highest energy
level and d electrons if they meet the criteria from the last
slide).
• How many valence electrons does Mg have?
• Electron configuration: [Ne]10 3s2
• Highest energy level is energy level 3 and there are only
two electrons there, which means that Mg has 2 valence
electrons.
Find valence electrons for the
following elements:
•
•
•
•
•
•
Be
O
F
Br
As
Fe
What are the types of bonding?
• Bonding is when two or more atoms come
together. Valence electrons are involved in
bonding.
• Ionic bonding is the bonding between metals and
non-metals. These compounds are solid at room
temperature, most of them dissolve in water and
they conduct electricity when they dissolve.
• Metallic bonding is the bonding between metals.
These compounds do not dissolve in water but
they do conduct electricity at room temperature.
Ionic Substance
• Covalent bonding-this is the bonding
between non-metals. Some of these
substances dissolve in water. Some do not,
but they never conduct electricity.
• Covalent Network solids-These are hard
solids made up of carbon or silicon atoms in
a network. An example of a covalent
network solid is …
Covalent Substance-Water
Covalent Network Solid
• What types of bonding do the following
substances exhibit?
NaOH
C(graphite)
Cr
CO
How do ions form?
• Ions form by losing or gaining electrons in
order to get a noble gas electron
configuration.
• Metals lose electrons and get a positive
charge. #electrons lost= +charge of metal
• Non-metals gain electrons and get a
negative charge. #electrons gained= charge of non-metal.
Ionic Bonding
• Ionic Bonding occurs between metals and
non-metals when electrons are transferred.
• In order to form ionic compounds, we must
first form ions.
• Ions are elements that have a charge
because they lose or gain valence electrons.
• Cations have a positive charge while anions
have a negative charge.
Ion formation
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Ions continued
• For metals, metals lose valence electrons until
they get a noble gas configuration and get a
positive charge (#electrons lost=positive charge of
element).
• Non-metals gain valence electrons until they get a
noble gas configuration and get a negative charge
(#electrons gained=negative charge of element).
• What is the charge of chlorine when it becomes an
ion? What about magnesium, sodium, aluminum?
How do we form ionic
compounds?
• We form ionic compounds by putting a metal ion
together with a non-metal ion. We cannot change
their charges but we can change the numbers of
each.
• We need to make ionic compounds that are
electrically neutral. In order to do this, we need to
add more of each ion until this occurs.
• A quick way to do this, if their charges do not
cancel, is to cross the coefficients (not the signs)
of the charges. Make sure that you have the
lowest whole number ratio between charges.
Making ionic compounds
• Make ionic compounds between:
Na and Cl
Al and F
Mg and N
What are polyatomic ions?
• Polyatomic ions are ions that have multiple
elements in them and they have an overall
charge. Look on your half sheet for those
charges. For polyatomic ions, place them in
parentheses before you cross the charges.
• Make ionic compounds with the following:
Na and SO4
Mg and NO3
K and PO4
Network of NaCl ions
• This shows how salt
crystals are put
together.
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How do we name ionic
compounds?
• Ionic compounds consist of two names that are based on
the elements that they contain. The first name is always
the metal that is in the compound (the element or
polyatomic ion that you write first in the compound). The
second name is for the non-metal in the compound.
• If the non-metal is a singular ion, then you use the name of
the element but you take off the last syllable and add -ide.
For example, chlorine becomes chloride when it is in an
ionic compound or oxygen becomes oxide.
• If the non-metal is a polyatomic ion then you use the name
of the polyatomic ion from your half sheet.
• For example, CaCl2 has the name of Calcium Chloride.
Questions
• Name the following ionic compounds:
NaCl
Cu(NO3)2
AlBr3
FeO
MgSO4
ZnS
• Convert the following back into ionic compounds:
Sodium Cyanide
Aluminum Chloride
Magnesium Oxide
Zinc Nitrate
How do we name covalent
compounds?
• For covalent compounds, you name both elements that are
involved in the compound. You write the name of the first
element and leave it unchanged. You then write the name
of the second element, after the first element, take off its
last syllable and add
-ide.
• You also must use prefixes with covalent compounds.
These prefixes denote how many atoms of each element
are in the compounds. The prefixes are: mono (1, use on
second element only), di (2), tri(3), tetra (4), penta (5),
hexa (6), hepta (7), octa (8), nona (9), and deca (10).
• For example, CO2 has the name of carbon dioxide.
• Name these covalent compounds:
CCl4
NH3
SO2
N2O4
• Write covalent compounds from these names:
Carbon Monoxide
Dihydrogen Monoxide
Dinitrogen Pentaoxide
Covalent Bonding
• Covalent bonding is between non-metals where
electrons are shared. Large biological molecules
have this type of bonding.
• In order to demonstrate covalent bonding, we first
must make Lewis dot symbols.
• You make these by taking an element’s valence
electrons and placing them around the element
symbol as dots. You place one around each of the
four corners and then add more as needed. Each
side can only have two electrons.
Covalent Bonding
• Unpaired electrons (1 dot) represent
available spots where covalent bonding can
occur.
• Paired up electrons represent lone pairs,
where bonding cannot occur.
• Make Lewis dot symbols for:
Ne S
F
N
C
Lewis structures
• Lewis structures are a 2-D representation of
covalent molecules.
• In order to make them, first split the molecule into
its component elements.
• Put Lewis dot symbols around each element.
• Pair up unpaired electrons, on different atoms, to
form covalent bonds (1 bond=2 shared electrons).
• Put it all together.
Do these examples:
• HF
• CH4
• PBr3
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What is the octet rule?
• The octet rule says that atoms tend to gain, lose or share
electrons so as to have eight electrons in their outer
electron shell. This means that all atoms, in a Lewis
structure, must have eight valence electrons around them
(they can be either bonded or lone pair electrons).
• Hydrogen and helium are exceptions to the octet rule.
There is one more element that is an exception to the octet
rule, which one is it?
• Some elements (like sulfur) can have more than eight
electrons around them, which means that they also break
the octet rule.
Practice
• Complete the following Lewis structures:
C2H5OH
CCl4
BH3
Multiple Bonds
• In order to form multiple bonds, you must first
make single bonds between the atoms. If there are
still unpaired electrons then you can start forming
double or triple bonds.
• Double bond=4 electrons between the atoms.
• Triple bond=6 electrons between the atoms.
• Make the following molecules:
O2
COHF
HCN
What do we need to know about
biochem?
• Carbon makes up the backbone of most biological
molecules. Carbon’s ability to form four bonds
(single, double, or even triple bonds) is
responsible for this.
• Large molecules (polymers), such as starch,
proteins, and nucleic acids are formed by
repetitive combinations of simple subunits.
• Amino acids are the subunits and hence building
blocks of proteins.
More Lewis Structures
• Draw Lewis Structures for the following:
N2
COH2
CO2
What is polarity?
• Polarity is the result of unequal sharing of
electrons. A polar covalent bond results when
there is a difference in electronegativity between
the elements involved in that bond.
• Due to this electronegativity difference, the
bonding electrons are shared unequally. They
usually spend more time around the more
electronegative atom. This creates a partial
positive and partial negative end of the atom.
• A dipole is created from this and it goes from
partial negative to partial positive end.
How do we know if a bond is
polar?
• We must compare the difference in electronegativities,
between the two elements, in order to determine whether
or not the bond is polar. The electronegativity for each
element is located on pg. 366.
• In order to do this, we must subtract the electronegativity
values for each element. We can classify the bond, based
on that difference, using the table on pg. 418. We can
classify the bond as non-polar, moderately polar, very
polar, and ionic.
• Classify the following bonds as non-polar, polar, or ionic:
• H-F
K and Cl
Br-Br
How do we know if a molecule is
polar?
• If a molecule is polar, there must be a net dipole.
• If a molecule is diatomic, with the same element,
then it is non-polar. If there are different elements
then it is polar.
• If a molecule contains only C-C or C-H bonds
then the molecule is non-polar.
• If a molecule contains a lone pair around the
central atom then that molecule is polar.
• If the central atom is carbon, and it is
bonded to the same element all around, then
the molecule is non-polar.
• If the central atom is carbon and it is
bonded to two or more different elements
all around then the molecule is polar.
• Polar or non-polar?
CH4
NH3
COHF
O2
Example of dipole on a polar
molecule.
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What are intermolecular forces?
• Intermolecular forces are forces between
molecules. Substances with strong intermolecular
forces will have a high boiling point and a low
freezing point.
• The types: Van Der Waal’s or London Dispersion
forces-these are the forces between non-polar
molecules. A temporary dipole forms between the
molecules. Due to their lack of true dipoles, nonpolar substances cannot interact with polar
substances, which is why they cannot dissolve in
polar substances such as water. These forces are
the weakest.
• Dipole-Dipole-These are the forces between polar
molecules. Their dipoles interact. The partial
positive ends interact with the partial negative
ends. These forces are middle strength.
• Hydrogen Bonding-These are between polar
substances that have a strong dipole, which forms
when hydrogen is bonded to N, O, or F. These are
the strongest behind ionic forces.
• Ionic forces-forces between ionic compounds.
These are very strong due to their charges.
Example of Hydrogen Bonding
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• Using intermolecular forces, which has the
highest boiling point between H2O, H2S,
CH4, and HCl?
• How do ammonia (NH3) molecules
interact? Can they dissolve in octane
(C8H18)?
What is a mole?
• Moles are a counting unit that we use in chemistry.
They allow us to convert between quantities.
• 1 mole=22.4 L at STP
• STP is standard temperature and pressure, which
is 0 C and 1 atmosphere of pressure.
• We do mole conversions using conversion factors
based on the equivalences. The units that you
want go on top and the units that you want to
cancel go on the bottom.
• Use your mole road map.
Not the chemistry mole
Problems
• How many moles are in 44.8 L of CO2 at
STP?
• How many liters are in 2 moles of O2 gas at
STP?
What is Avogadro’s number?
• Amadeo Avogadro, an Italian scientist,
hypothesized that equal volumes of gas, at equal
temperatures and pressures, contain the same
number of particles. Scientists eventually proved
him right.
• Avogadro’s number is 6.02 x 1023 particles. These
particles could be atoms (singular elements),
molecules (covalent compounds), or formula units
(ionic compounds).
• 1 mole= 6.02 x 1023 particles (atoms, formula
units, molecules)
Avogadro’s hypothesis
Problems-use mole road map
• How many moles are in 5 x 1020 molecules
of CO2?
• How many atoms of Cu are in 2 mol of Cu?
What is the gram formula mass?
• The gram formula mass (gfm) of a substance is defined as
the sum of all the atomic masses of the elements in a
compound multiplied by the number of each respective
element.
• To find the gfm, find what elements and how many of each
element you have in a compound. Multiply the number of
each element by its atomic mass and add them up.
• For example, in CaCl2, you have 1 Ca and 2 Cl. That
corresponds to the following gfm:
40.08 (atomic mass of 1 Ca) + 2 x 35.5 (atomic mass of Cl x
2)=111.08 g
Find the gfm for the following
substances:
•
•
•
•
NaCl
Cu(NO3)2
MgSO4
CO2
What about moles and mass?
• 1 mole = gfm of that element or compound.
• This is the most important mole relationship
and we will use it all year long.
• How many moles are in 50 g of NaCl?
• How many grams are in 2 moles of
Zn(OH)2?
• How many moles are in 20 g of MnO2?
Mole Review
• Remember that:
1 mole= 6.02 x 1023 particles
1 mole = 22.4 L of a gas at STP
1 mole= gfm of a substance
Problems
• How many moles are in 1.204 x 1024
formula units of NaCl?
• How many liters are in 3 moles of O2 at
STP?
• How many moles are in 100 g of Pb(NO3)2?
Mole Review Transparency
• Convert the following to moles:
20 g NaCl
3 x 1024 atoms He
40 g AlPO4
10 g Ba(OH)2
44.6 L N2
4 x 1012 atoms Ca
Convert the following from moles
2 mol CaSO4 to grams 3 mol He to atoms
5.6 mol N2 to L
2.3 mol Al(NO3)3 to grams
6 mol Zn to atoms
5 mol NaOH to grams
What are two-step mole
conversions?
• Two-step mole conversions allow us to convert
between two different units on the mole road map.
They involve two conversions. The first
conversion takes you from your initial quantity to
moles. The second conversion takes you from
moles back out to your wanted quantity.
• In these two-step conversions, moles are the
intermediate.
• This is a very powerful tool. For example, it can
tell us the mass of one molecule of carbon dioxide.
Problems
• How many grams are in 11.2 L of Ne gas at
STP?
• How many grams are in 1 molecule of
Carbon Dioxide (CO2)?
What is percent composition?
• Percent composition is the percent by mass of
each element in a compound. The percent mass of
an element in a compound is the number of grams
of the element (that make up the gfm of the
compound) divided by the total gfm of the
compound, multiplied by 100%.
• %mass of element E= (grams of element in one
mole of the compound)/gfm of compound) x 100
• That formula applies for all elements in the
compound.
Problems
• Calculate the percent composition of each
element in C3H8.
• Calculate the percent composition of each
element in NaNO3.
What is the empirical formula?
• The empirical formula of a compound gives
the lowest whole number ratio of the atoms
of the elements in a compound.
• The empirical formula is valuable because it
tells us the kinds and relative count (atoms
or moles of each element) in molecules or
formula units of a compound.
How do we calculate empirical
formulas?
• We calculate empirical formulas by using the percent composition of
each element in a compound and turning it into grams. For example,
68% N becomes 68 g N. We do this for each element.
• We then convert the grams of each element into moles using the
appropriate conversion and gfm.
• Then, we find the lowest number of moles and divide all mole
quantities (for all elements) by it in order to get the lowest whole
numbers.
• If you get a number that is close to a whole number when you divide
by the lowest number of moles then round it up to that number. For
example, you would round 2.9 to 3. For the number 2.5, you would
not round it. You would multiply all mole numbers by 2 so that the 2.5
became 5.
• Write the empirical formula using these lowest whole numbers of
moles that you found. The number of moles goes next to each element
in the compound.
Problem
• Determine the empirical formula of a
compound that is 25.9% nitrogen and
74.1% oxygen.
• Determine the empirical formula of a
compound that is 16.49% N, 56.53% O, and
26.97% Na.
What are molecular formulas?
• The molecular formula is the same as its experimentally determined
empirical formula or it is some whole number multiple of it.
• To calculate molecular formulas, you first must find the empirical
formula of a compound.
• Then you must find the gfm of that empirical formula.
• Divide the gfm of the molecular formula (usually given to you and
called the gram molecular mass) by the gfm of the empirical formula.
This should give you a whole number to multiply by.
• Multiply the empirical formula by that whole number (that affects the
number of each element in the compound) in order to get the molecular
formula.
Problems
• Ethane has a gram molecular mass of 30
g/mol. If ethane’s percent composition is
79.9% carbon and 20.1% hydrogen, what is
its molecular formula?
What are chemical reactions?
• Chemical reactions (sometimes called chemical
equations) describe chemical processes. In
chemical reactions, an arrow separates the
formulas of the reactants (on the left) from the
formulas of the products (on the right)
• We read reactions from left to right (like a book).
• Chemical reactions occur inside of us and in all
living things. They help us break down our food
intake, send signals throughout our body, and
eliminate harmful waste products or toxins.
Chemical Reaction (burning of
wood)
What information can we get
from chemical reactions?
• We can get the reactants and products of a
chemical reaction. We can also get their states.
The states of a reactant or product are written next
to that reactant or product in subscript parentheses.
(s) means solid, (l) means liquid, (g) means gas,
and (aq) means aqueous or dissolved in water.
• Above the reaction arrow, you might see a
chemical formula. That is a catalyst, which helps
the reaction go but is not part of the reaction itself.
You also could see H, which means that you had
to add heat.
Problems
• Describe the following reactions in words:
• CaCO3 (s) CaO (s) + CO2 (g)
• S (s) + O2 (g) SO2 (g)
What is conservation of mass?
• Conservation of mass states that you can neither
create nor destroy mass. You can only change its
form. Basically, it states that you cannot create
something out of nothing. For example, you
cannot turn lead into gold.
• At the subatomic particle level, (quantum
mechanics) we are seeing subatomic particles
appear seemingly out of nowhere. This research
will give us a better understanding of where these
particles came from during and after the initial
expansion of the universe (Big Bang)
Why is conservation of mass
important?
• Conservation of mass influences how we
write chemical equations. According to
conservation of mass, all chemical reactions
must be balanced. A balanced chemical
reaction has the same number of atoms of
each element on both sides of the reaction.
How do we balance reactions?
• We balance reactions by taking an inventory of
each element on both sides of the reaction.
• You cannot change the formulas of the reactants or
products, you can only change the coefficients out
in front of each reactant or product.
• Remember that these coefficients affect all
elements in the compound and they represent
moles.
Example of balancing reaction
• Balance this reaction:
• H2 + O2 H2O
• There are two hydrogen on the reactant side and two oxygen on the
reactant side.
• There are two hydrogen and one oxygen on the product side. This
means that the oxygen is unequal. We would balance that by putting a
2 in front of the water.
• H2 + O2  2H2O
• This balances our oxygen but not hydrogen. We would then have to
put what in front of the hydrogen on the reactant side?
• Our final equation would look like:
• 2H2 + O2 2H2O
Practice
• Balance these reactions
• C2H6 + O2  CO2 + H2O
• NaCl + Al(NO3)3  NaNO3 + AlCl3
What are the types of reactions?
• There are five types of reactions (know them because I
may ask you about them on a quiz…)
• Combination reactions occur when two or more elements
or compounds combine to form a different compound. A +
B AB
• Decomposition reactions occur when a compound decays
into its component elements or new compounds. AB  A
+B
• Single replacement reactions occur when a metal reacts
with an ionic compound. This metal replaces the metal
that is in the ionic compound. This forms a new
compound and places the other metal on its own. A + BC
B + AC where A and B are metals.
• Double replacement reactions occur when two
ionic compounds react. The metals, in these ionic
compounds, switch places and form two new ionic
compounds.
AB + CD  AD + CB
• Combustion reactions occur when elements
(usually hydrocarbons) burn with oxygen (you
must provide an initial jump of energy, like heat,
to start them) to produce carbon dioxide and
water.
CnHn + O2 CO2 + H2O
Double replacement reaction
Lead Iodide is the solid
Combustion in a diesel engine
Balance the following reactions
and identify their type
• C3H8 + O2  CO2 + H2O
• H2O2 H2O + O2
• Mg + HCl  MgCl2 + H2
What are single replacement
reactions?
• In a single replacement reaction, one metal
reacts with an ionic compound. The metal,
that is on its own, replaces the metal that is
in the ionic compound. As a result, a new
compound forms and a different metal is on
its own.
How do we write single
replacement reactions?
• First, you have the metals switch places (if there is only
one metal, then the non-metals switch places). Write each
element (or polyatomic ion) in its singular form on the
product side of the equation. For example:
Mg + AlCl3  MgCl + Al
• Balance the charges on the new ionic compound that you
have formed. Once you have done that, then balance the
reaction overall.
• Thus it should be:
Mg + AlCl3  MgCl2 + Al (balance charges)
3Mg + 2AlCl3  3MgCl2 + 2Al (balance entire reaction)
Practice
• Write the following reactions:
• Mg + Zn(NO3)2 
• Al + H2SO4 
What are double replacement
reactions?
• In a double replacement reaction, two ionic
compounds react. The metals switch places
and form two new ionic compounds.
• As a result, two new ionic compounds form.
How do we write double
replacement reactions?
• First, the metals (in both ionic compounds) switch places.
Write each element (or polyatomic ion) in its singular form
on the product side of the equation. For example:
BaCl2 + K2CO3  BaCO3 + KCl
• Balance the charges on the new ionic compounds that you
have formed. Once you have done that, then balance the
reaction overall.
BaCl2 + K2CO3  BaCO3 + KCl (charges balanced)
BaCl2 + K2CO3 BaCO3 + 2KCl (balance entire reaction)
Practice
• Write the following double replacement
reaction:
• MgBr2 + Zn(NO3)2 
What are combustion reactions?
• In combustion reactions, an element or a
compound (usually a hydrocarbon) reacts
with oxygen producing heat and light.
• Combustion reactions need an initial burst
of energy (through a spark) to get going.
• Combustion reactions occur when
hydrocarbons burn in oxygen to produce
carbon dioxide and water.
How do we write combustion
reactions?
• In a combustion reaction, the hydrocarbon burns
in oxygen to produce carbon dioxide and water.
The products of this type of reaction are always
carbon dioxide and water.
• Thus we would write:
• C3H8 + O2  CO2 + H2O
• Then balance the overall reaction to get:
• C3H8 + 5O2  3CO2 + 4H2O
Practice
• Write out and balance the following
combustion reactions:
• C5H10 + O2 
• C6H12O6 + O2 
Combustion Reaction
What is the activity series of
metals?
• The activity series of metals (pg. 191 in your
textbook) is a list of metals in order of decreasing
reactivity.
• The activity series determines whether or not a
single replacement reaction will occur. Single
replacement reactions will only occur if the metal,
that is on its own, is more active than the metal
that is in the ionic compound.
• If the metal that is on its own is not more active
than the metal that is in the ionic compound then
no reaction will occur.
• Remember that the activity series of metals
only applies to single replacement reactions.
• Write out the following single replacement
reactions (write no reaction if a reaction will
not occur):
• Ag + NaNO3 
• Mg + FeCl3 
What are net ionic reactions?
• You can write single or double replacement reactions as a
complete ionic reaction, which is a reaction that shows
dissolved ionic compounds as their component ions. You
can only split up ionic compounds into their component
ions if they have (aq) next to them. You cannot do this for
compounds that have (l), (g), or (s).
• Spectator ions are ions that are not directly involved in the
reaction. These ions appear on both sides of the reaction.
You can cancel them out.
• After doing this, you can write a net ionic equation that
indicates only the particles that actually take part in the
reaction.
Practice
• Write a net ionic equation for the following:
AgNO3 (aq) + NaCl (aq) AgCl(s)+ NaNO3 (aq)
• You can split the three aqueous compounds into their
component ions like so:
Ag+ (aq) + NO3- (aq) + Na+(aq) + Cl- (aq) AgCl(s)+ Na+(aq) +
NO3- (aq)
• Thus, the spectator ions are Na+ and NO3- because they
appear on both sides of the reaction unchanged (still in
aqueous form). You can cross these out and write the net
ionic equation as:
Ag+ (aq) + Cl- (aq) AgCl(s)
Practice
• Write out the following as a net ionic
equation:
• Pb(s) + 2AgNO3 (aq) Pb(NO3)2 (aq) + 2Ag (s)
What is temperature?
• Temperature is the measure of the average speed
of molecules in a substance. Heat raises the
temperature because it causes the molecules to
move faster.
• Temperature is measured in three scales:
Farenheit, Celsius, and Kelvin. Remember that
water freezes at 0 C and boils at 100 C. In the
Kelvin scale, there are no negative temperatures.
• How does temperature affect you? How do you
respond to high temperatures? What about low
temperatures?
Thermometer
How do we convert between
temperature scales?
• There are formulas that allow us to convert
between temperature scales.
• C=5/9 (F-32)
• F=9/5(C) + 32
• K=C + 273
What is absolute zero?
• In the Kelvin scale, there are no negative
temperatures. 0 Kelvin is called absolute
zero because it is the lowest possible
temperature. At this temperature, all
molecular motion stops, gases occupy no
volume, and have no pressure.
Temperature Conversions
• Complete the following temperature
conversions:
• Convert 15 C to Farenheit and Kelvin
• Convert 98.6 F to Celsius and Kelvin
• Convert 373 K to Celsius and Farenheit
Temperature Activity
• Convert the following temperatures to Kelvin:
• 273 C
50 C
• 27 C
45 C
• 450 C
100 C
75 C
90 C
730 C
• Convert the following temperatures to Celsius:
• 273 K
0K
50 K
• 345 K
400 K
760 K
• 540 K
335 K
600 K
What is pressure?
• Pressure is the force per unit area. Gases
move from areas of high pressure to areas
of low pressure.
• We measure pressure in units of
atmospheres (atm), millimeters of mercury
(mmHg), Kilopascales (Kpa), and pounds
per square inch (psi).
Pressure Defined
Weight of air on you due to
pressure
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
What are mole ratios?
• Mole ratios are a comparison of the number of
moles of two different substances (either reactants
or products) in a reaction. Mole ratios can
compare two reactants, two products, or a reactant
and a product.
• To find mole ratios, simply look at the coefficients
in front of the two substances that you are looking
at. Place those two numbers on either side of a
colon (:)
• For example, in the following reaction:
• 3NaNO3 +AlCl33NaCl + Al(NO3)3
• The mole ratio between NaNO3 and
Al(NO3)3 is 3 mol NaNO3: 1 mol Al(NO3)3.
• What is the mole ratio between AlCl3 and
NaCl? What about between NaNO3 and
AlCl3?
What is Stoichiometry?
• Stoichiometry studies the quantitative
aspects of chemical reactions.
Stoichiometry allows us to predict how
much reactant we will need or how much
product we will form using those quantities
of reactants.
What are mole to mole
conversions?
• In Stoichiometry, we must learn a new conversion
called mole to mole conversions. These
conversions come from the mole ratios (of
coefficients) in a reaction. These conversions
allow us to convert from moles of one substance
to moles of another substance.
• As always, the moles that you want go on top and
the moles that you are trying to cancel go on the
bottom.
Example
• 3NaNO3 +AlCl33NaCl + Al(NO3)3
• For the above reaction, how many moles of NaCl can form from 5
moles of AlCl3?
• First we must write our given quantity outside of the conversion factor.
• Then we must find the mole ratio between NaCl and AlCl3 and it is:
• Then we must set up our conversion factor based on the mole
quantities from this reaction. The moles that we want (NaCl) go on
top and the moles that we want to cancel (AlCl3) go on the bottom of
the conversion factor. The numbers of each come from the coefficients
in the reaction. You can then convert the numbers.
• 3NaNO3 +AlCl33NaCl + Al(NO3)3
• Using the above reaction, how many moles
of NaNO3 are needed to react with 2 moles
AlCl3? How many moles of Al(NO3)3 can
form from 2 moles AlCl3?
What are mass to mass
conversions?
• Mass to mass conversions allow us to convert from grams
of one substance to grams of another substance. Mole to
mole conversions are still used to convert between moles
of two different substances.
• If the problem gives you a mass of one reactant or product,
you must first convert it into moles.
• Then, convert it into moles of a different substance using
the mole to mole conversion.
• Then, convert from moles of that new substance back into
grams of that new substance. In all, you will need up to
three conversion factors to do this.
How many conversion factors
will you need?
• You will need 1 conversion factor to go from
moles of one substance to moles of another
substance.
• You will need 2 conversion factors to go from
grams of one substance to moles of another
substance or from moles of one substance to
grams of another substance.
• You will need 3 conversion factors to go from
grams of one substance to grams of another
substance.
Practice
• 3NaNO3 +AlCl33NaCl + Al(NO3)3
• How many grams of NaCl can form from 3
moles of AlCl3?
• How many grams of Al(NO3)3 can form
from 30 grams of AlCl3?
More Practice
• 2Al(NO3)3 +3Zn 3Zn(NO3)2 + 2Al
• How many moles of Al can form from 3
moles of Zn?
• How many grams of Zn(NO3)2 can form
from 60 grams of Zn?
What is a limiting reactant?
• The limiting reactant is the reactant that is
completely used up in a chemical reaction. None
of it remains after a reaction becomes complete. It
all goes to form products, which is why the
limiting reactant determines how much product
can form. The limiting reactant is the reactant that
forms the least amount of product.
• The theoretical yield is how much product
(usually in grams) should form in a chemical
reaction. The theoretical yield is the amount of
product that the limiting reactant forms.
The limiting reactant determines
how much gas will form
How do we determine the
limiting reactant?
• The limiting reactant is the reactant that forms the
least amount of product. You must convert the
quantities of reactant to the same product. You
can only convert these quantities to the same
product because you need to compare them. The
reactant that forms the least amount of product is
limiting.
• Also, the quantity of product that the limiting
reactant forms is the theoretical yield.
Practice
• 2Al(NO3)3 +3Zn 3Zn(NO3)2 + 2Al
• If 20 grams Al(NO3)3 and 50 grams of Zn react,
which reactant is limiting? What is the theoretical
yield for Al?
• If 100 grams Al(NO3)3 and 75 grams of Zn react,
which reactant is limiting? What is the theoretical
yield for Al?
• Mg + 2HCl MgCl2 + H2
• Given the above reaction, if 15 grams Mg
and 40 grams HCl react, which reactant is
limiting? How much Hydrogen (H2) should
form?
What is the excess reactant?
• The excess reactant is the reactant that is
left over at the end of a reaction. It is the
reactant that forms the larger amount of
product when you do limiting reactant
calculations.
• You can determine how much of the excess
reactant is left over at the end of the
reaction.
• You can tell how much excess reactant is left over by first
finding the limiting reactant.
• You then can identify which reactant is in excess.
• You then take your given quantity of limiting reactant and
convert it to grams of the excess reactant in order to see
how many grams of excess reactant react with the limiting
reactant.
• You then subtract the given quantity of excess reactant by
how much excess reactant reacts with the limiting reactant.
This tells you how much excess reactant is left over at the
end of the reaction.
Practice
• Mg + 2HCl MgCl2 + H2
• Given the above reaction, if 15 grams Mg
and 40 grams HCl react, which reactant is in
excess? How much of this reactant is left at
the end of the reaction?
What is the percent yield?
• Percent yield compares how much product
you made (actual yield) to how much
product you should have made (theoretical
yield). A percent yield above 90% is very
good.
• Percent yield=(actual yield/theoretical
yield) x 100
Practice
• Mg + 2HCl MgCl2 + H2
• Given the above reaction, if 1 gram of H2
forms, what is the percent yield?
Saturation Example
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
Example of how pressure affects
gas solutes.
• Increasing the pressure
above the solution
makes more gas
molecules dissolve in
a solvent.
• What companies
would use this
principle?
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
What is concentration?
• Concentration is the amount of solute that is
dissolved in a given amount of solvent. Solutions
are concentrated if they have a large amount of
solute dissolved in a given amount of solvent.
• Why would we care about the concentration of a
substance?
• All substances are toxic to humans in certain
concentrations.
Concentrated Detergent
What do we measure
concentration in?
• We measure concentration using four
different measurements:
• Molarity (M)=moles solute/liters solvent
• Grams per Liter (g/L)=grams solute/liters
solvent
• Which is easier to determine?
Practice
• I dissolve 60 grams of Pb(NO3)2 into .5 L of
water. What is the concentration of this
mixture in molarity? What about grams per
Liter?
• I dissolve 40 grams of NaOH into 2 L of
water. Express this concentration in
molarity and grams per Liter.
What are the other concentration
units?
• Parts per million (ppm)=
milligrams of solute/kilograms of solvent
• Percent Composition (%)=
(grams solute/mL of solvent)x100
Practice Problems
• A sample of river water has 4 mg of lead in .2 Kg
of water. Express this concentration in ppm.
• A fertilizer says that it has “5% NH4NO3 by
volume.” You test this batch of fertilizer and find
that it has 10 mg NH4NO3 per 500 mL bottle. Is
this bottle’s label true in terms of how much
NH4NO3 it actually has?
What is heat?
• Heat is a form of energy that flows from higher
temperatures (hot) to cold (lower temperatures).
Heat is released or absorbed during any
mechanical work.
• You cannot feel heat, you only feel heat transfer.
• Exothermic-heat is released. Endothermic-heat is
absorbed.
• Why do you feel cold on a winter day?
Image of heat transfer
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
What is a phase change?
• The phases of matter are solid, liquid, and
gas. A phase change occurs when
substances change their state of matter.
This is a physical change.
• During a phase change, the temperature of a
substance does not change. All of the
absorbed energy is used to make the
substance change state.
What are the phase changes?
• Solid to liquid: Melting
• Liquid to Solid: Freezing
• Liquid to gas:
Boiling/Evaporation
• Gas to liquid:
Condensation
• Solid to gas: Sublimation
(see picture)
• Gas to Solid: Solidication.
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
• System- This is the specific part of the universe
that you are focusing on.
• Surroundings- This is everything outside of the
system.
• Endothermic reaction-This is a process that
absorbs heat from the surroundings.
• Exothermic reaction-This is a process that releases
heat to the surroundings.
What is specific heat?
• Specific heat is the amount of heat needed
to heat 1 g of a substance by 1 C. Every
substance has its own unique specific heat.
• We can use this formula if we have specific
heat: H=cm T, where H equals the heat
change in joules, c=specific heat, m=mass,
and T=Tfinal-Tinitial.
Heat calculations
• How much heat is needed to raise the
temperature of water from 30 C to 50 C?
(Specific heat of water=4.18 J/g C
• How much heat is released when 10 g
copper cools from 100 C to 50 C? Specific
heat of copper=.33 J/g C
How do we use ∆H=cm∆T for
two substances?
• In this situation, one substance loses heat (the one at the
higher temperature) and one substance absorbs that same
amount of heat (the one at lower temperature). Since they
absorb or release the same amount of heat, we can relate
them. Also, the final temperature of the system is the same
for both substances.
• We use this equation:
-cm∆T =
cm∆T
Heat released heat absorbed
Where ∆T=Tfinal - Tinitial
• We use this formula to find specific heats of objects among
other things.
• I drop a 20 g piece of copper, at 100 C, into 20 g
of water at 25 C. What is the final temperature of
my system? Cwater=4.18 J/g C and Ccopper=.387 J/g
C
• I drop 20 grams of an unknown substance, at 90C,
into 50 grams of water at 20 C. If the final
temperature of the water and unknown substance
are 24 C, what is the specific heat of the uknown
substance?
How do we calculate heat for
phase changes?
• For phase changes, we use the following
formulas:
• H=m Hfus
H=m Hvap
• Hfus=heat it takes to melt or freeze
something.
• Hvap=heat it takes to boil or condense
something.
Image of Ice Melting
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
Heat calculations
• How much heat does it take to melt 20
grams of ice at 0 C? (Heat of fusion=333
J/g)
• I add 400 J of heat to 40 g of ice. How
much of this ice melts? How much of it is
left over?
How can we use the additive
properties of heat?
• We can use the additive properties of heat in
problems that involve both phase and temperature
changes (for example, a problem that goes from
ice at one temperature to water at a different
temperature). When a temperature change occurs,
use cm∆T but use either m∆Hfus or m∆Hvap for
phase changes.
• Add up the heats that you calculate as you change
the temperature.
Problems
• How much heat is needed to melt 50 g ice,
at -10 C into water at 15 C?
• How much heat is needed to change 20 g of
water, at 50 C, into gas at 110C?
What is Pressure?
• Pressure is a force per unit area. Pressure is the
force per unit area. Gases move from areas of
high pressure to areas of low pressure.
• We measure pressure in units of atmospheres
(atm), millimeters of mercury (mmHg),
Kilopascales (Kpa), and pounds per square inch
(psi).
• We can convert between pressure units.
How do we convert between
pressure units?
• Like any conversion, the units you want go
on top of the conversion and the units that
you want to cancel go on the bottom of the
conversion.
• Use these equivalencies:
1 atm= 760 mmHg = 101.3 kPa
Problems
• Convert 5 atm into kilopascales.
• What represents more pressure, .25 atm or
550 mmHg?
What is Charles’ Law?
• Charles’ law is the volume-temperature
relationship. It is a direct relationship and
pressure and amount of gas are kept constant.
• V1/T1=V2/T2
• V1-initial volume, T1=initial temperature, V2-final
volume, T2-final temperature.
• Remember that temperatures must be in Kelvin!!
Charles’ Law Example
Calculations
• I have a gas at 30 K and 4 L. What is my
new volume when my temperature
increases to 60 K?
• I have a gas at 25 C and 2 L. What is my
new temperature if the volume decreases
to .5 L?
What is Boyle’s Law?
• Boyle’s law is the pressure-volume relationship. It
is an indirect relationship and temperature and
amount of gas are kept constant.
• P1V1=P2V2
• V1-initial volume, P1=initial pressure, V2-final
volume, P2-final pressure.
Too Much Pressure…
Calculations
• I have a gas at 2 atm and 3 L. What is
my new pressure if my volume increases
to 6 L?
• I have a gas at 400 kPa and 7 L. What is
my new volume if the pressure increases
to 500 kPa?
What is Gay-Lussac’s Law?
• Gay Lussac’s law is the pressure-temperature
relationship. It is a direct relationship and volume
and amount of gas are kept constant.
• P1/T1=P2/T2
• P1-initial pressure, T1=initial temperature, P2-final
pressure, T2-final temperature.
• Remember that temperatures must be in Kelvin!!
Example
Calculations
• I have a gas at 600 mmHg and 30 C.
What is my new temperature if my
pressure increases to 800 mmHg?
• I have a gas at 400 mmHg and 20 C.
What is my new pressure if my
temperature increases to 100 C?
What gas laws have we learned
so far?
• Charles: V1/T1=V2/T2 (Pressure is constant)
• Boyles: P1V1=P2V2 (Temperature is
constant)
• Gay-Lussac’s: P1/T1=P2/T2 (Volume is
constant)
• Remember to use Kelvin and that the
amount of gas is constant in all of these.
• I have a gas at 3 L and 5 atm. What is my
new pressure if the volume decreases to .5
L?
• I have a tire at a pressure of 32 psi and a
temperature of 30 C. What is my new
temperature if the pressure becomes 45 psi?
What is the Combined Gas Law?
• The combined gas law is a combination of the
other three laws that we have learned. You can
derive the other three gas laws by keeping one of
the variables, in the combined gas law, constant.
• P1V1/T1=P2V2/T2
• Only the amount of gas is kept constant. P1-initial
pressure, V1-initial volume, T1=initial temperature,
P2-final pressure, V2-final volume, T2-final
temperature.
Calculations
• A sample of air occupies a volume of 3 L
at 50 C and 4 atm. What is the volume of
air at 100 C and 8 atm?
• A sample of air occupies a volume of 4 L
at 30 C and 2 atm. What is the new
pressure of this gas at 2 L and 25 C?
What are acids? What are bases?
• Acids are hydrogen donors while bases are
hydrogen acceptors.
• Acids are corrosive and taste sour. Bases are
bitter, slippery, and caustic. Both can cause
serious chemical burns.
• Strong acids completely dissociate in water while
weak acids only partially dissociate in water. The
same is true for strong bases vs. weak bases.
• What are some examples of acids or bases?
Strong vs. Weak acids
• These images show
the difference between
strong and weak acids.
• Does the same thing
hold true for strong
bases vs. weak bases?
Quic kTime™ and a
TIFF (Unc ompres sed) dec ompres sor
are needed to see this pic ture.
Quic kTime™ and a
TIFF (Unc ompres sed) dec ompres sor
are needed to see this pic ture.
Strong Acids can be dangerous
What is the pH scale?
• The pH scale is a scale that goes from 0 to 14.
• The pH scale allows us to measure acidity and
basicity. We classify substances according to their
pH. pH is a measure of the hydrogen ion
concentration.
pH=-log[H+]
pOH=-log[OH-]
pH+pOH=14
Here is the pH scale
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pH classifications
• Please label the following classifications on your
pH scale:
• pH 0-3=strong acid
• pH 4-6=weak acid
• pH6-8=neutral
• pH 9-11=weak base
• pH 11-14=strong base
• If something has a pH of 5, what do we classify it
as?
What are reaction rates?
• Reaction rates measure how fast a reaction
occurs. They are proportional to the
decrease in reactant concentration over time
or the increase in product concentration
over time.
• Consider this reaction: AB
• Rate= -[A]/ t= [B]/ t
What affects reaction rates?
• In order for a reaction to occur, the reactants must have
enough energy to overcome the activation energy barrier
(see board). The reaction, even if it is spontaneous, will
not occur unless it can overcome the activation energy
barrier.
• Temperature: This affects reaction rate because
temperature measures the speed of molecules. If the
temperature is high then the molecules are moving faster
and you will have more collisions with higher energy. This
will increase the reaction rate. At a lower temperature, the
molecules move slower and you will have less high energy
collisions, which will slow down the reaction rate.
• Pressure affects reaction rates in the same way that
temperature does.
• Concentration-this is the amount of each reactant.
If you increase the concentrations of the reactants,
there are more reactants in the vessel. This
increases the probability of high energy collisions
and increases the reaction rate. Decreasing the
concentrations of reactants decreases the reaction
rate because there are less reactants and less high
energy collisions.
• Surface Area- increasing the surface area
increases the reaction rate because there are
more high energy collisions between
reactants.
What is a catalyst?
• A catalyst is something that increases the
reaction rate by lowering the activation
energy. A catalyst can be a metal ion, or a
huge protein. The biological catalysts are
enzymes.
Example of a Catalyst
• This is a picture of an
enzyme in your body.
• You can see where it
binds the reactants so
that they can react.
This binding lowers
the activation energy.
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What is Equilibrium?
• Equilibrium reactions are represented by a
double arrow. For example: AB. You
cannot apply normal stoichiometric
calculations to these reactions because there
is both a forward and reverse reaction that
can occur.
• Equilibrium occurs when the forward and
reverse reaction rates are equal.
What is Le-Chatlier’s Principle?
• Le-Chatlier’s principle states that if a stress
is applied to a system at equilibrium then
the system changes to relieve that stress.
• This shift will either be toward the product
side or toward the reactant side depending
on the stress that is put on the equilibrium.
This only applies to aqueous and gaseous
reactants or products.
What affects equilibrium?
• Concentration-Increasing the concentration of a
product pushes the equilibrium reaction back
towards the reactant side. Increasing the
concentration of a reactant pushes the equilibrium
reaction in the direction of the product side.
• Decreasing the concentration of a product pulls
the equilibrium towards the product side.
Decreasing the concentration of a reactant pulls
the equilibrium towards the reactant side.
Temperature
• Increasing the temperature of an equilibrium
causes the reaction to shift towards the side that is
endothermic (absorbs heat) or does not have the
heat term.
• Decreasing the temperature of an equilibrium
causes the reaction to shift towards the side that is
exothermic (releases heat) or the side that has the
heat term.
Pressure
• Pressure only affects gaseous reactants or products
(g) not aqueous reactants or products. It only
affects equilibrium systems with an unequal
number of moles of gaseous reactants and
products.
• Increasing the pressure forces the equilibrium to
shift towards the side with fewer moles of gas.
• Decreasing the pressure forces the equilibrium to
shift towards the side with more moles of gas.
Problem
• Consider the following equilibrium reaction:
N2(g) + 3H2(g) 2NH3(g) + heat
Which way does the equilibrium shift if I make the
following changes to the system:
Increase N2 Decrease H2 concentration
Decrease NH3 concentration
Increase temperature
Decrease temperature
Increase pressure
Example of Le-Chatlier’s
Principle
• Ammomia Synthesis
(the Haber process):
N2(g) + 3H2(g) + heat
2NH3(g)
• How can I maximize
the synthesis of
ammonia using LeChatlier’s principle?
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Atomic Structure Review
• What did we learn about in our atomic structure
unit?
• Atoms-definition, number of particles in each,
isotopes, isotope notation.
• Periodic table families and trends.
• Nuclear Chemistry-types of nuclear decay, fission,
fusion, and writing nuclear decay reactions.
• Electron configurations-long-hand and short-hand
along with finding the number of valence
electrons.
Problems
• How many protons, neutrons, and electrons are in
carbon 14?
• Which has the higher electronegativity between C,
N, and O?
• Write out the nuclear decay reaction that occurs
when U-235 undergoes alpha decay.
• What is the electron configuration for bromine and
how many valence electrons does it have?
What did we learn about
bonding?
• We learned about the types of bonding-ionic,
metallic, covalent, and covalent network.
• We learned about how ions form and how to make
ionic compounds.
• We learned how to name compounds.
• We learned how to make Lewis structures to
represent covalent structures.
• We learned about some biochem topics also.
• What type of bonding do we find in CO2?
• Form an ionic compound with Mg and Cl,
and Mg and NO3.
• Name those compounds from second bullet.
• Draw the structure of CH4, N2H4.
• What are the subunits of proteins?
What did we learn about moles
and reactions?
• We learned how to convert moles to mass,
volume at STP, and particles. We learned
two step conversions, gfm, empirical
formulae, and molecular formulae.
• We learned conservation of mass and how
to balance reactions.
• We learned how to write reactions.
Problems
• Convert 22 grams of Mg(NO3)2 to moles.
• Convert 20 L of oxygen gas to moles
• Convert 3x1024 atoms of nitrogen into
grams of nitrogen.
• Balance the following reaction:
• NaCl + Al(NO3)3  NaNO3 + AlCl3
What did we learn about
Stoichiometry?
• We learned how to do mole to mole conversions,
mass to mass conversions, calculate limiting
reactants, and how to calculate excess reactants.
• 2Al(NO3)3 +3Zn 3Zn(NO3)2 + 2Al
• If 20 grams Al(NO3)3 and 50 grams of Zn react,
which reactant is limiting? What is the theoretical
yield for Al?
•
What did we learn about
solutions?
• We learned about how substances dissolve, what
makes them dissolve, how we can get more solute
to dissolve into a solvent, what a solute and
solvent are, and how to calculate concentration.
• How can I get more solute to dissolve into a
solvent?
• What is the molarity of a solution where I dissolve
20 g of MgCl2 into 2L of water?
What did we learn about acids/bases,
reaction rates, and equilibrium?
• The definition of an acid/base, the difference
between strong acids/bases and weak acids/bases,
the pH scale and classifications.
• For reaction rates, we learned the definition and
what we can do to speed up a reaction rate.
• For equilibrium, we learned the definition and
how to apply Le-Chatlier’s principle.
• What is an acid? How would we classify a solution that has a pH of 9?
• How can we increase the rate of a reaction?
• Consider the following equilibrium reaction:
N2(g) + 3H2(g) + heat 2NH3(g)
Which way does the equilibrium shift if I make the following changes to
the system:
Increase N2
Decrease H2 concentration
Decrease NH3 concentration
Increase temperature
Decrease temperature
Increase pressure
What did we learn about
thermodynamics?
• We learned about the types of phase changes, the
definition of heat, how to calculate heat during a
temperature change, or phase change, or both.
• How much heat is needed to raise 20 g of water’s
temperature from 30 C to 50 C? (Specific heat of
water=4.18 J/g C
• How much heat is needed to melt 20 grams of ice?
Heat of fusion for ice=333 J/g
What did we learn about the gas
laws?
• We learned the definition of temperature, pressure,
absolute zero, and how to use the four gas laws.
• A sample of air occupies a volume of 3 L at 50
C and 4 atm. What is the volume of air at 100
C and 8 atm?
• I have a gas at 2 atm and 3 L. What is my new
pressure if my volume increases to 6 L?
Warm-up
• Why would we want to know how much energy it
takes to melt ice? Why is this important to us?
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Exit Slip
• What improvements could you make to the
experiment in order to improve your results?
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What is Hess’ Law?
• Hess’s Law of Heat Summation states that if you
add two or more thermochemical equations to give
a final equation, then you can also add the heat
changes to give the final heat change.
• In order to use Hess’s Law, you must cancel out
substances or elements (if they are on their own)
that are on both sides of the reaction when you add
the reactions.
• When you use Hess’s Law, make sure you identify
which reaction has the reactants or products that
you want.
• If the product you want is written as a reactant,
then you can switch the reaction around. If you do
this, the sign on the heat (∆H) changes (it can go
from a negative to a positive).
• If you need to multiply a reaction (in order to
decrease or increase the amount of a product or
reactant) by a constant, then you must multiply
∆H by that same constant.
• Once you have the needed reaction, add up the
heats.
Solve a Hess’s Law problem
• Solve this Hess’s law problem:
Find ∆H for the following reaction:
2Cu(s) + O2(g)2CuO(s)
Given:
CuO(s) + Cu(s) Cu2O(s) ∆H=-11.3 KJ
Cu2O(s) + 1/2O2(g)2CuO(s) ∆H=-114.6 kJ
Which reactions contain the reactants you want?
Which ones have the products that you want?
What is the Ideal Gas Law?
•
The ideal gas law relates temperature, pressure,
and volume to amount of gas in moles.
• This law forces us to make ideal gas
assumptions which are:
a. There are no attractive forces between gas
molecules.
b. Gas molecules occupy no volume.
• Real gases do not behave this way all the time,
but the ideal gas law gives us a good
approximation.
• The ideal gas law is this formula:
PV=nRT
P=pressure in atm, V=volume in L, n=moles
of gas, R=.0821, T=Temperature in Kelvin.
• You can make changes to this equation in
order to make it work better for real gases.
Example
Problems
• I have a gas at 2 atm of pressure, 2L of volume, and 30 C.
How many moles of this gas do I have?
• 50 grams of Carbon dioxide gas is at a temperature of 150
C, and a volume of 3 L. What is the pressure of this gas?
• I have oxygen gas in a container at 3 atm, 5 L, and 25 C.
How many grams of oxygen gas are in this container?
What is Dalton’s Law of Partial
Pressures?
• Partial Pressure is the pressure contribution that
each gas, in a mixture, makes to the total pressure
of the gas mixture.
• Dalton’s law of partial pressure states that at
constant volume and temperature, the total
pressure exerted by a mixture of gases is equal to
the sum of the partial pressures of the component
gases.
• Ptotal=P1 + P2 + P3 +…
Example
Problem
• Air contains oxygen (O2), nitrogen (N2),
carbon dioxide (CO2) and trace amounts of
other gases. What is the partial pressure of
oxygen P(O2) at 101.3 kPa of pressure if
P(N2)=79.1 kPa, P(CO2)=.04 kPA, and
P(other gases)=.94 kPA?
What are oxidation and reduction
reactions?
• Oxidation and Reduction reactions (called
redox) occur when electrons are transferred
between reactants.
What is oxidation?
• Oxidation occurs when an organic
compound adds oxygen or decreases its
number of hydrogens.
• We also define it as the loss of electrons.
• Substances that oxidize will donate
electrons, which means that they are a
reducing agent.
What is reduction?
• Reduction is when organic compounds lose
oxygen or gain hydrogens.
• We also define it as the gain of electrons.
• The substance that is reduced in an
oxidation-reduction reaction is called the
oxidizing agent.
Redox Reaction Example
What are the rules for assigning
oxidation numbers?
1.
2.
3.
4.
The oxidation number of a monoatomic ion is equal to
its ionic charge. (Sodium has a +1 charge)
The oxidation number of hydrogen in a compound is
always +1 except in metal hydrides such as NaH (where
it is -1)
The oxidation number of oxygen in a compound is
always -2.
The oxidation number of any uncombined element is
zero. For example, the oxidation number of solid
potassium metal (K) is zero. What about the oxygen in
O2?
5. For any neutral compound, the sum of the
oxidation numbers of the atoms in the compound
must equal zero.
6. For a polyatomic ion, the sum of the oxidation
numbers must equal the ionic charge of the ion.
• What is the oxidation number of each element in
the following compounds?
a. SO2
b. CO3-2
c. K2SO4
How do we identify oxidation
and reduction in reactions?
• An increase in oxidation number signifies
oxidation.
• A decrease in oxidation number signifies
reduction.
• In order to determine which element is oxidized or
reduced we must assign oxidation numbers to each
element, on both sides of the reaction.
• Then we determine which element had a decrease
in its oxidation number (reduced) and which
element had an increase in its oxidation number
(oxidation).
• Use the information on the previous slide to
solve the following problem:
Use the changes in oxidation numbers to
identify which elements are oxidized and
reduced in this reaction:
Zn(s) + 2MnO2(s) + 2NH4Cl(aq)
ZnCl2(aq) + Mn2O3(s) + 2NH3(g) + H2O(l)
How do we balance redox
reactions?
• We use the oxidation number change method to
balance redox reactions by comparing the
increases and decreases in oxidation numbers.
• This principle works because the total number of
electrons gained in reduction must equal the total
number of electrons lost in oxidation.
• In order to balance redox reactions we must use
the following steps:
1.
2.
3.
4.
5.
Assign oxidation numbers to all the atoms in the
equation.
Identify which atoms are oxidized and which are
reduced.
Use a line to connect the atoms that undergo oxidation
and those that undergo reduction. Write the oxidation
number change at the midpoint of each line.
Make the total increase in oxidation number equal to the
total decrease in oxidation number by using the
appropriate coefficients.
Finally, check to be sure that the equation is balanced for
both atoms and charge.
Use the method highlighted
above to solve the following
problem:
• Balance this redox reaction by using the
oxidation number change method.
K2Cr2O7(aq) + H2O(l) + S(s)KOH(aq) +
Cr2O3(aq) + SO2(g)
How do we use half reactions to
balance Redox reactions?
• Half-reactions are equations that show
either the reduction or the oxidation of a
species in an oxidation-reduction reaction.
• The half-reaction method is used to balance
redox reactions by balancing the oxidation
and reduction half reactions.
• This method is particularly useful when you
are balancing ionic reactions.
What are the steps to balancing
redox reactions using half reactions?
1.
2.
3.
4.
5.
Write the equation in ionic form.
Identify the oxidized species and the reduced species.
Write out their half reactions.
Balance each half reaction in terms of charge. Make
sure that the charges are equal on both sides of the half
reaction.
Multiply each half reaction by an appropriate number to
make the electron charges equal. This will allow you to
cancel out electrons in each half reaction.
Combine half-reactions. Check to make sure all
numbers are balanced.
Use the method outlined above
for the following reaction:
• Balance this reaction using the half-reaction
method:
MnO4-(aq) + Cl-(g)Mn+2(aq) + Cl2(g)
But this reaction does not
balance!!
• The above reaction is not balanced for the oxygen atoms.
How can we balance it?
• Under acidic conditions, we do the following:
• Add oxygen to the either the reactant or product side using
H2O. Add as many waters as you need to balance out the
oxygen. (If you have four oxygens on the reactant side,
you will need four waters on the product side).
Continuing under Acidic
conditions:
• Add H+ to the opposite side so that we can balance
out the hydrogens (if you added four waters to one
side, then you would add eight H+ ions to the other
side) that we just added to the other side of the
reaction when we added water.
• Check if charges balance.
What about under basic
conditions?
• Under Basic conditions, we balance the
reaction under acidic conditions then we
add OH- to both sides of the equation. The
number of hydroxides (OH-) added equals
the number of H+ ions.
• Together, the H+ and OH- ions form water.
Use these new waters to cancel out the
waters on the other side of the reaction.
Now back to our problem:
• Balance this reaction under acidic and basic
conditions:
MnO4-(aq) + Cl-(g)Mn+2(aq) + Cl2(g)
Electrochemistry
• Electrochemical cells convert chemical
energy into electrical energy and vice versa.
All electrochemical processes involve redox
reactions.
• An electric current is produced from the
flow of electrons between the ions in the
redox reaction.
Volataic Cells
• Voltaic cells are electrochemical cells that are used to
convert chemical energy into electrical energy. The energy
is produced by spontaneous redox reactions within the cell.
• In order to create a voltaic cell, you must split the reaction
into two half reactions. A half-cell is one part of a voltaic
cell in which oxidation or reduction occurs. The two halfcells are separated by a porous partition.
• A salt bridge is a tube that contains a conducting solution.
A salt bride allows the passage of ions from one half cell to
the other it prevents the two solutions from mixing
completely.
How do voltaic cells work?
• The half-cell (or electrode) where oxidation occurs
is called the anode.
• The half-cell where reduction occurs is called the
cathode.
• Electrons flow from anode to cathode. If you
connect the two half-cells via an external wire
(and with a salt bridge) you can generate an
electrical current, which is how batteries work.
Here is an example of an
electrochemical cell:
How do we calculate cell
potentials?
• In order to calculate cell potentials, we must look
at the standard reduction potential.
• For half reactions, the standard reduction potential
is always written showing the species undergoing
reduction. Electrons are reactants.
• For standard reduction potential you must have
standard conditions (25 C, and 1 atm of pressure).
• The standard reduction potential for reducing
H+(aq) to H2(g) is defined to be 0 V
• In order to calculate standard potential (E) for any
electrochemical cell we must use this formula:
• E= Eanode+ Ecathode
• The anode is the reaction with the more negative half-cell
reduction potential.
• The cathode is the reaction with the more positive half-cell
reduction potential.
• For the anode, you will change the sign on its voltage
because that voltage is for reduction. To change the
voltage to oxidation, you must change the sign on the E
for that reaction.
• For any electrochemcial cell to be spontaneous
(occur without the input of energy) its standard
cell potential must be positive.
• What is the standard cell potential (E) for the
following redox reaction and is this reaction
spontaneous?
Zn(s) + Cu+2(aq)Zn+2(aq) + Cu(s)
Make sure you identify the oxidation and
reduction reactions. Look up E for each half
reaction (remember to change sign on voltage for
anode) in your textbook.
What are alkanes?
• Alkanes are hydrocarbons that contain only
single covalent bonds. They can be in a
long and straight chain molecule or they can
be branched.
• We name the straight-chain alkane based on
its number of carbons. We use the
following prefixes and add -ane to the end:
What are functional groups?
• Functional groups are a specific arrangement of
atoms in an organic compound that is capable of
characteristic chemical reactions.
• When we represent the general form of functional
groups, we use an -R group to represent carbon
chains that are attached to the functional group.
• The table on pg. 712 (AW book) shows the
general compound forms and their functional
groups. This table should be on your note page
for the upcoming organic chemistry quiz…
Prefixes based on # of carbons in
a straight chain:
1 carbon-meth
2 carbons-eth
3 carbons-prop
4 carbons-but
5 carbons-pent
6 carbons-hex
7 carbons-hep
8 carbons-oct
9 carbons-non
10 carbons-dec
Images of alkanes
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How do we name alkanes?
1.
2.
3.
4.
5.
Find the longest continuous chain of carbons in the
molecule. If it is four carbons, then you call it butane (as
the parent structure)
Number the carbons in the main chain in sequence. This
will allow you to identify each functional group’s
position on the chain.
Add numbers to the names of the substituent or
functional groups to identify their positions on the chain.
Use prefixes to indicate the appearance of a group more
than once in the structure.
List the names of the functional groups in alphabetical
order.
How do we name the functional
groups?
• If the functional group is a hydrocarbon then it is
called an alkyl group. You name them by
removing the -ane end and adding -yl to the end.
Make sure you number where they are on the
chain when you name the compound.
• For halogens, the are named by taking off -ine at
the end and adding -o (chlorine becomes chloro).
Make sure you number where they are on the
chain when you name the compound.
• Halocarbons are synthesized and used as
anesthetics and insecticides.
How do we name alkenes and
alkynes?
• Alkenes are organic compounds containing
carbon-carbon double bonds.
• Alkynes are organic compounds containing
carbon-carbon triple bonds.
• Organic compounds that contain double and triple
bonds are called unsaturated compounds because
they contain fewer than the maximum number of
hydrogens in their structure.
• Alkanes that contain the maximum number of
hydrogens are called saturated compounds (no
carbon-carbon multiple bonds).
• To name alkenes or alkynes, follow the same
naming rules and numbering. However, the
continuous chain carbon must contain the double
or triple bond. The ending changes from -ane to ene for alkenes or -yne for alkynes.
• In naming these compounds, you must number the
chain so that the carbon atoms of the double or
triple bond gets the lowest possible numbers.
Make sure to put the carbon number where the
double bond is on the alkene or alkyne when you
name the compound.
Name the compounds on the
board
What are the types of organic
reactions?
• Substitution: The replacement of an atom
or group of atoms by another atom or group
of atoms. For example, a halogen can
replace the hydrogen on an alkane to
produce a halocarbon.
• Let’s do some examples on the board.
Many reactions we see involve
double or triple bonds
• Addition reaction-this occurs when a substance is added at
the double or triple bond of an alkene or alkyne. Addition
reactions allow us to introduce new functional groups into
organic molecules. For example, you can add two
chlorines on the carbons that make up the double bond.
• If you add water to an alkene then it is called a hydration
reaction.
• Hydrogenation reaction-this occurs when you add
hydrogen (H2) to a carbon-carbon double bond and
produce an alkane. This reaction requires a catalyst in the
form of platinum (Pt) or palladium (Pd)
What about oxidation/reduction?
• Dehydrogenation reactions occur when you lose hydrogen
and create a carbon-carbon double or triple bond. Strong
heating and a catalyst are necessary to do this. This causes
oxidation of the molecule (loss of hydrogen is considered
oxidation).
• You can oxidize molecules by adding more oxygens.
Alkanes can become alcohols if you oxidize them. These
alcohols can become aldehydes or ketones after oxidation
depending on their functional groups. These can then be
oxidized into esters or alcohols depending on their
functional groups.
• Hydrogenation reactions are reduction (adding
more hydrogen)
What is hydrolization?
• Hydrolization of esters occur if an ester is
heated with water for several hours in
strong acid or base solutions. This produces
a carboxylic acid and an alcohol.
What is Biochemistry?
• Biochemistry covers large macromolecules that
are found in most life forms.
• These macromolecules (large molecules) form
from polymers, which are continuous repetitions
of the same smaller molecules (monomers).
• The types of macromolecules are: carbohydrates,
lipids (fats), proteins, and nucleic acids (DNA).
What are carbohydrates?
• Carbohydrates are monomers and polymers of aldehydes
and ketones that have numerous hydroxy groups attached.
They have the formula of Cn(H2O)n.
• The simplest carbohydrates are simple sugars called
monosaccharides such as glucose and fructose (fruit
sugar).
• If two sugars come together via a polysaccharide bond,
then we call them disaccharides. The most famous is a
combination of fructose and glucose, which creates
sucrose. This is the sugar that we use in candy and to
sweeten our drinks.
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What are polysaccharides?
• The linkage of many monosaccharide monomers produces
polysaccharides.
• Examples include:
Starch-it is made up of glucose molecules and serves to store
glucose molecules in plants. Starch can partially dissolve
in water and is edible.
Glycogen-this is the animal form of starch. Humans store
glucose in this way
Cellulose-it is made up of glucose. The orientation of the
bond linking the glucose monomers is different from the
orientation in starch. As a result, cellulose is insoluble and
can only be digested by a limited number of organisms.
Continuation of cellulose
• Cellulose forms a rigid structure with other
cellulose molecules. Cellulose is used as a
structural polysaccharide. Plant cell walls,
as in wood, are made of cellulose.
• Humans cannot digest cellulose. Cellulose
is then considered what type of nutrient?
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Lipids
• Lipids are a large class of relatively water-insoluble
compounds that include fats, oils, and waxes. They can
dissolve in organic solvents such as ether.
• Triglycerides are triesters of long-chain fatty acids and
glycerol. They are the major component of animal fats and
oils. Triglycerides are simple lipids and they are important
as the storage form of fat in the human body.
• Triglycerides are essentially esters that we can hydrolize
by boiling them with aqueous sodium hydroxide. This
makes soap.
What are phospholipids?
• Phospholipids are a special type of lipid that
contain a polar head (phosphate group) and
a non-polar tail (hydrocarbons).
• These unique lipids make up the cell
membrane in animal cells. It is called the
phospholipid bilayer.
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Image of a phospholipid bilayer
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•
•
The polar head faces the outside and inside of the cell so that it can interact
with molecules outside or inside of the cell.
The non-polar tails face toward the inside on both layers so that they can
interact with one another. This makes the cell membrane relatively
impermeable because molecules have a hard time getting through the nonpolar tails. Large molecules cannot fit between the phospholipids.
Here are some images of Lipids
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What is the difference between
saturated vs. unsaturated fats?
• Saturated fats have no carbon-carbon double bonds in their
structure. This allows them to interact well (VDW forces)
and melt at higher temperatures.
• Unsaturated fats have carbon-carbon double bonds in their
structures. This puts a kink in the fatty acid chains, which
means that they do not interact as well (VDW forces) and
melt at lower temperatures. They are usually liquid at
room temperature.
• What household products are made up of mostly saturated
fat? Which ones are made up of most unsaturated fats?
Unsaturated vs. Saturated Fats
• Here are two images
that show the
difference between
saturated and
unsaturated fats.
QuickTime™ and a
TIF F (Uncompressed) decompressor
are needed to see this picture.
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
What are amino acids?
• Amino acids are compounds that contain amino (NH2) and carboxylic acid (-COOH) groups in the
same molecule.
• There are 20 amino acids, all with the
aforementioned functional groups. They also have
another functional group and hydrogen around the
central carbon.
• At human blood pH (7.4), these amino acids exist
in what is called their zwitterion form.
Here is an image of amino acid
and zwitterion form (know it)
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
What are peptides?
• Amino acids are the building blocks of proteins.
They come together via peptide bonds, which
form between one carboxylic acid of one amino
acid and the amino group of the next amino acid.
• Peptides are any combination of amino acids that
have come together via peptide bonds.
• Polypeptides are peptides with more than 10
amino acids.
• Proteins are peptides with more than about 100
amino acids.
Here is an image of a peptide
bond
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
How do proteins get their various
shapes?
• Proteins fold into various shapes known as alpha
helices or beta sheets. They may fold into more
complex shapes based on intermolecular forces,
covalent bonds formed between cysteine
molecules, and electrostatic attractions.
• Proteins can be denatured by changes in
temperature and pH. These changes disrupt the
attractions that keep them folded in the correct
conformation.
Images of protein shapes
• Here is an image of an
alpha helix.
• Here is an image of a
beta sheet.
QuickTi me™ and a
TIFF ( Uncompressed) decompressor
are needed to see thi s pi ctur e.
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
What are Enzymes?
• Enzymes are proteins that act as biological
catalysts. Enzymes are unchanged during a
chemical reaction. Enzymes help necessary body
reactions to occur.
• The enzyme binds substrates (the molecules on
which an enzyme acts) at its active site. This is
where the enzyme converts substrates into
products. The enzyme holds the substrates in
place (through intermolecular forces) so that they
can react or it even reacts with the substrate. Each
enzyme is specific to a certain substrate.
What affects enzyme function?
• Some enzymes will not function without a coenzyme.
These are metal ions or small organic molecules that need
to be present in order for an enzyme catalyzed reaction to
occur. Many coenzymes are nutrients that we take in
everyday such as zinc, iron, and ascorbic acid (vitamin C).
• Enzymes can only function in certain temperature or pH
ranges. They break down and do not function if they are
not within that range. What body conditions (due to
sickness) can impair enzyme function?
• Some poisons inhibit enzyme function. They attach to the
enzyme at the active site and prevent the substrate from
binding. They can also attach elsewhere on the enzyme
and change its shape so that it cannot bind the substrate.
Here is an enzyme:
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
What are nucleic acids?
• Nucleic acids are polymers found primarily in cell nuclei.
The most common are DNA (deoxyribonucleic acid) or
RNA (ribonucleic acid). DNA stores the information
needed to make proteins and governs the reproduction and
growth of cells and new organisms. RNA transmits the
information stored in DNA.
• Nucleotides are the subunits of DNA and RNA. They
contain a phosphate group, a five carbon sugar (ribose),
and a nitrogen base. For DNA its ribose is missing an
oxygen. The four bases are adenine (A), Thymine (T),
Cytosine (C), and Guanine (G).
How does DNA come together?
• DNA winds in a double helix shape where the negative
charges on the phosphate groups repel each other (this
keeps the helices at the correct distance apart). In the
middle, the bases interact through hydrogen bonding.
• The bases pair in this way:A-T and C-G. RNA contains
these bases but it does not have Thymine, it contains Uracil
(U), which pairs up with adenine.
• Find the corresponding DNA sequence:
TTCGATTGGCAAA
Here is the DNA double helix:
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
Final Exam Review
• The following items will be on our semester
2 final:
• Stoichiometry, solutions, acids and bases,
reaction rates, equilibrium, thermodynamics
(including Hess’s law), gas laws (including
Ideal Gas Law), oxidation/reduction,
organic chemistry, and biochemistry.
Review Problems
• Find the oxidation numbers on the elements in the
following compounds:
CaSO4
MgO
O2
• In the reaction below, determine the oxidized
element, reduced element, oxidizing agent, and
reducing agent.
Mg + CaSO4 MgSO4 + Ca
• What is the pressure of 3 mol of gas at a volume
of 3 L and a temperature of 30 C?
O. Chem + Biochem Problems
• What is the complimentary DNA sequence to
AATCCG?
• What molecule makes up a cell membrane? What
does it form?
• What are the subunits of proteins? What is the
zwitterion form of this molecule? What is an
enzyme? What affects its activity?
• What are the polysaccharides?
• Name the compounds on the board (use IUPAC
rules) and write out the reactions that are there.