I wasn't particularly happy with my grades as they dropped my overall gpa to a 2.63 (need a 2.7 for nursing) I've one semester left before I can get in so hopefully my grades in there will affect it for a positive response.
My grades this semester were...
Chemistry - C
Microbiology - B
Microbiology Lab - B
Sociology - C
Psychology - B
Thursday, May 21, 2009
Thursday, January 29, 2009
Chapter 3 Notes for Chemistry Exam
3.1 – The Periodic Law and Table:
Scientists were trying to find order or some sort of pattern to explain the chemistry of
the elements.
Periodic Law – When arranged in increasing atomic number, elements with similar
chemical properties occur at regular (periodic) intervals.
Dmitri Mendeleev was the first to observe this periodic behavior and he organized
the elements into a table with vertical columns or groups and horizontal rows or
periods.
Group or Family – A vertical column of elements in the periodic table having
similar chemical properties. Scientists are still fighting over how to label the groups,
but we will number them from left to right.
Period – A horizontal row of elements across the periodic table. Periods are
numbered from top to bottom.
Each element in the periodic table belongs to both a group and a period.
3.2 – Electronic Arrangements in Atoms:
So why do elements have similar properties at regular periodic intervals?
The answer lies in their electrons.
Model A Model B
Ernest Rutherford proposed a planetary model for the atom where e-s orbit the
nucleus of an atom like planets orbit the sun.
Niels Bohr proposed that e-s can only orbit the nucleus at fixed distances from the
nucleus, and therefore e-s can only have specific energies.
Bohr also proposed that e-s can only jump to orbits of different distances from the
nucleus when energy is absorbed or released.
Absorbing energy causes an e- to jump farther away from nucleus to higher energy
and less stable orbits.
Releasing energy causes the e- to fall back
down closer to the nucleus to lower energy
and more stable orbits.
Erwin Schrodinger further refined this
model stating that the precise paths of e- can’t
be determined accurately (Like Bohr
thought). According to Schrodinger, the
location and energy of electrons around the
nucleus can be specified using 3 terms: shell,
subshell, and orbital. This model is called
the quantum mechanical model.
Shell – A location and energy of e-s around a
nucleus that is designated by a value for n,
where n = 1, 2, 3.
The lowest energy shell is assigned n = 1, and the next lowest is n = 2, etc.
The higher the value of n, the farther away from the nucleus the e-s are and the more
energy they have.
Subshell – A sub-compartment of a shell designated by the letters: s, p, d, f.
Subshells are identified using shell # (value for n) and the subshell letter (s, p, d, f).
Example: for n = 3, there are 3 subshells designated 3s, 3p, and 3d
All e-s within a specific subshell have the same energy (they are the same distance
from the nucleus).
Subshells are further divided into 1 or more atomic orbitals.
Orbital – A specific volume of space around the nucleus of an atom where e-s of the
same energy move. All the orbitals of a subshell have the same value of n.
The volume of space around the nucleus of each subshell is different.
All e-s in the same orbital have the same energy, regardless of the orbital they are in.
Example: An e- in a 4d orbital has the same energy as any other e- in another one of
the five 4d orbitals.
All s subshells consist of 1 singular orbital.
All p subshells consist of 3 orbitals
All d subshells consist of 5 orbitals
All f subshells consist of 7 orbitals
Each orbital within a subshell can contain a
maximum of 2 e-.
Energy Diagram of Atomic Shells, Subshell, and Orbitals
Practice Problems:
What is the e- configuration for Nitrogen?
What is the e- configuration for Calcium?
What is the e- configuration for Nickel?
What is the e- configuration for Chlorine?
3.3 The Shell Model and Chemical Properties:
The arrangement of e-s into orbitals, subshells, and shells provides and explanation
for the similarities in chemical properties of various elements.
All the elements in a specific group (column) have the same number of electrons in
their outer shell.
This shell is called the valence shell.
Valence Shell – The outermost shell of an element that contains the highest energy
electrons in that atom.
Elements with the same number of e-s in their valence shell display similar properties.
The number of e-s in a valence shell is identical the Roman numeral above the group
on the periodic table.
The n value for the valence shell increases by one as you go down the periodic table.
Practice Problem:
How many valence e-s does oxygen have?
How many valence e-s does magnesium have?
How many valence e-s does Yttrium need to lose or gain to obtain a full valence?
How many valence e-s does Sulfur need to lose or gain to obtain a full valence?
How many valence e-s does Fluorine need to lose or gain to obtain a full valence?
3.4 Electronic Configurations:
Electronic Configurations – The detailed arrangement of e-s indicated by a specific
notation. (1s2, 2s2, 2p6, 3s2, 3p6, 3d10)
Hund’s Rule – Electrons will always fill any available empty orbitals of the same
energy before pairing up to share an orbital. Think about people sitting on a bus.
Pauli Exclusion Principle – Only e- spinning in opposite directions can
simultaneously occupy the same orbital.
Order for filling atomic orbitals with e-s.
Helpful hint for determining e- configuration:
Noble Gas Configuration – An electronic configuration consisting of completely
filled s and p outermost subshells. 8 valence e-s, very stable, essentially inert. Far
Right Group of elements are called the “Noble gases”, they do not form bonds with
any other atoms. They are completely inert. Given the choice, atoms always want to
gain or lose e-s to obtain a full octet of 8 valence e-s.
3.5 Another Look at the Periodic Table:
Elements in the same family have the same # of valence e-s. (F, Cl, Br, I, & At each
have 7). As a result they all have similar chemical properties.
Distinguishing Electron – the last and highest-energy electron in an element.
Distinguishing e-s are categorized by the type of subshell (s, p, d, or f) they are found
in. (See figure 3.9)
Noble gases – Elements in the far right side of the periodic table. All are gases at
room temp and all are inert. They all have completely filled valence orbitals.
The elements can also be classified as: metals, nonmetals, and metalloids.
Metals – Found in the left two thirds of the periodic table. They have the following
properties:
• High thermal conductivity – transmit heat well
• High electrical conductivity – transmit electricity well
• Ductility – can be stretched in wires
• Malleability – can be hammered into thins sheets
• Metallic luster – shiny, “metallic” appearance
Nonmetals – Found in the right one third of the PT. Often brittle, powdery solids or
gases and have properties opposite those of metals.
Metalloids – Have properties between those of metals and nonmetals. They include:
B, Si, Ge, As, Sb, Te, At. They form a “staircase” separating metals from nonmetals
on the PT. (See figure 3.12)
3.6 Properties and Trends within the Periodic Table:
Elements become less metallic from left to right across a period and from bottom to
top in a group.
Elements decrease in size from left to right across a period and increase in size from
top to bottom down a group.
Each added proton results in a stronger attractive force between the nucleus and the
electrons.
So, each added proton pulls the e- cloud closer to the nucleus, shrinking the atom as
you move from left to right across the PT.
Scale Drawings and Atomic Radii of Various
Atoms Across the Periodic Table
Ionization energy – the energy
required to remove an electron from a
neutral atom.
Na Na+ + e-
The higher the value of the ionization
energy, the harder it is to remove an e-
.
Ionization energy increases from left
to right across a period and decreases
in size from top to bottom down a
group.
As the size of the atom increases, the
e-s are farther away from the positive,
attractive forces in the nucleus, so they are held more loosely than e-s closer to the
nucleus and therefore easier to remove.
Note: When sodium loses 1 e-, the remaining valence e-s have a full octet and the
sodium cation has the same e- configuration as Neon (a noble gas). The Na+ cation is
very stable and forms many commonly used salts, such as NaCl (table salt) and
NaHCO3 (baking soda).
The tendency of metals to lose their e-s (and conduct electricity) is because by giving
a way a couple of e-s to an atom that is a few e-s short of its own octet, the metal can
have a completely filled valence and increase its stability.
This tendency to lose e-s results in metals holding their e-s less tightly, allowing e-s
(and therefore electricity) to flow easily from one atom of the metal to the other. It’s
almost like hot potato, no one atom wants the e-s, because it would rather lose it and
have a full octet, so it passes it down the line to the next atom and BOOM, electricity.
Electronegativity (EN) – the attraction an element has for e-s, the higher the EN, the
more e-s are pulled towards the atom.
Electronegativity increases from left to right across a period and decreases from top
to bottom down a group. Fluorine is the most electronegative atom. Francium is the
least electronegative.
Scientists were trying to find order or some sort of pattern to explain the chemistry of
the elements.
Periodic Law – When arranged in increasing atomic number, elements with similar
chemical properties occur at regular (periodic) intervals.
Dmitri Mendeleev was the first to observe this periodic behavior and he organized
the elements into a table with vertical columns or groups and horizontal rows or
periods.
Group or Family – A vertical column of elements in the periodic table having
similar chemical properties. Scientists are still fighting over how to label the groups,
but we will number them from left to right.
Period – A horizontal row of elements across the periodic table. Periods are
numbered from top to bottom.
Each element in the periodic table belongs to both a group and a period.
3.2 – Electronic Arrangements in Atoms:
So why do elements have similar properties at regular periodic intervals?
The answer lies in their electrons.
Model A Model B
Ernest Rutherford proposed a planetary model for the atom where e-s orbit the
nucleus of an atom like planets orbit the sun.
Niels Bohr proposed that e-s can only orbit the nucleus at fixed distances from the
nucleus, and therefore e-s can only have specific energies.
Bohr also proposed that e-s can only jump to orbits of different distances from the
nucleus when energy is absorbed or released.
Absorbing energy causes an e- to jump farther away from nucleus to higher energy
and less stable orbits.
Releasing energy causes the e- to fall back
down closer to the nucleus to lower energy
and more stable orbits.
Erwin Schrodinger further refined this
model stating that the precise paths of e- can’t
be determined accurately (Like Bohr
thought). According to Schrodinger, the
location and energy of electrons around the
nucleus can be specified using 3 terms: shell,
subshell, and orbital. This model is called
the quantum mechanical model.
Shell – A location and energy of e-s around a
nucleus that is designated by a value for n,
where n = 1, 2, 3.
The lowest energy shell is assigned n = 1, and the next lowest is n = 2, etc.
The higher the value of n, the farther away from the nucleus the e-s are and the more
energy they have.
Subshell – A sub-compartment of a shell designated by the letters: s, p, d, f.
Subshells are identified using shell # (value for n) and the subshell letter (s, p, d, f).
Example: for n = 3, there are 3 subshells designated 3s, 3p, and 3d
All e-s within a specific subshell have the same energy (they are the same distance
from the nucleus).
Subshells are further divided into 1 or more atomic orbitals.
Orbital – A specific volume of space around the nucleus of an atom where e-s of the
same energy move. All the orbitals of a subshell have the same value of n.
The volume of space around the nucleus of each subshell is different.
All e-s in the same orbital have the same energy, regardless of the orbital they are in.
Example: An e- in a 4d orbital has the same energy as any other e- in another one of
the five 4d orbitals.
All s subshells consist of 1 singular orbital.
All p subshells consist of 3 orbitals
All d subshells consist of 5 orbitals
All f subshells consist of 7 orbitals
Each orbital within a subshell can contain a
maximum of 2 e-.
Energy Diagram of Atomic Shells, Subshell, and Orbitals
Practice Problems:
What is the e- configuration for Nitrogen?
What is the e- configuration for Calcium?
What is the e- configuration for Nickel?
What is the e- configuration for Chlorine?
3.3 The Shell Model and Chemical Properties:
The arrangement of e-s into orbitals, subshells, and shells provides and explanation
for the similarities in chemical properties of various elements.
All the elements in a specific group (column) have the same number of electrons in
their outer shell.
This shell is called the valence shell.
Valence Shell – The outermost shell of an element that contains the highest energy
electrons in that atom.
Elements with the same number of e-s in their valence shell display similar properties.
The number of e-s in a valence shell is identical the Roman numeral above the group
on the periodic table.
The n value for the valence shell increases by one as you go down the periodic table.
Practice Problem:
How many valence e-s does oxygen have?
How many valence e-s does magnesium have?
How many valence e-s does Yttrium need to lose or gain to obtain a full valence?
How many valence e-s does Sulfur need to lose or gain to obtain a full valence?
How many valence e-s does Fluorine need to lose or gain to obtain a full valence?
3.4 Electronic Configurations:
Electronic Configurations – The detailed arrangement of e-s indicated by a specific
notation. (1s2, 2s2, 2p6, 3s2, 3p6, 3d10)
Hund’s Rule – Electrons will always fill any available empty orbitals of the same
energy before pairing up to share an orbital. Think about people sitting on a bus.
Pauli Exclusion Principle – Only e- spinning in opposite directions can
simultaneously occupy the same orbital.
Order for filling atomic orbitals with e-s.
Helpful hint for determining e- configuration:
Noble Gas Configuration – An electronic configuration consisting of completely
filled s and p outermost subshells. 8 valence e-s, very stable, essentially inert. Far
Right Group of elements are called the “Noble gases”, they do not form bonds with
any other atoms. They are completely inert. Given the choice, atoms always want to
gain or lose e-s to obtain a full octet of 8 valence e-s.
3.5 Another Look at the Periodic Table:
Elements in the same family have the same # of valence e-s. (F, Cl, Br, I, & At each
have 7). As a result they all have similar chemical properties.
Distinguishing Electron – the last and highest-energy electron in an element.
Distinguishing e-s are categorized by the type of subshell (s, p, d, or f) they are found
in. (See figure 3.9)
Noble gases – Elements in the far right side of the periodic table. All are gases at
room temp and all are inert. They all have completely filled valence orbitals.
The elements can also be classified as: metals, nonmetals, and metalloids.
Metals – Found in the left two thirds of the periodic table. They have the following
properties:
• High thermal conductivity – transmit heat well
• High electrical conductivity – transmit electricity well
• Ductility – can be stretched in wires
• Malleability – can be hammered into thins sheets
• Metallic luster – shiny, “metallic” appearance
Nonmetals – Found in the right one third of the PT. Often brittle, powdery solids or
gases and have properties opposite those of metals.
Metalloids – Have properties between those of metals and nonmetals. They include:
B, Si, Ge, As, Sb, Te, At. They form a “staircase” separating metals from nonmetals
on the PT. (See figure 3.12)
3.6 Properties and Trends within the Periodic Table:
Elements become less metallic from left to right across a period and from bottom to
top in a group.
Elements decrease in size from left to right across a period and increase in size from
top to bottom down a group.
Each added proton results in a stronger attractive force between the nucleus and the
electrons.
So, each added proton pulls the e- cloud closer to the nucleus, shrinking the atom as
you move from left to right across the PT.
Scale Drawings and Atomic Radii of Various
Atoms Across the Periodic Table
Ionization energy – the energy
required to remove an electron from a
neutral atom.
Na Na+ + e-
The higher the value of the ionization
energy, the harder it is to remove an e-
.
Ionization energy increases from left
to right across a period and decreases
in size from top to bottom down a
group.
As the size of the atom increases, the
e-s are farther away from the positive,
attractive forces in the nucleus, so they are held more loosely than e-s closer to the
nucleus and therefore easier to remove.
Note: When sodium loses 1 e-, the remaining valence e-s have a full octet and the
sodium cation has the same e- configuration as Neon (a noble gas). The Na+ cation is
very stable and forms many commonly used salts, such as NaCl (table salt) and
NaHCO3 (baking soda).
The tendency of metals to lose their e-s (and conduct electricity) is because by giving
a way a couple of e-s to an atom that is a few e-s short of its own octet, the metal can
have a completely filled valence and increase its stability.
This tendency to lose e-s results in metals holding their e-s less tightly, allowing e-s
(and therefore electricity) to flow easily from one atom of the metal to the other. It’s
almost like hot potato, no one atom wants the e-s, because it would rather lose it and
have a full octet, so it passes it down the line to the next atom and BOOM, electricity.
Electronegativity (EN) – the attraction an element has for e-s, the higher the EN, the
more e-s are pulled towards the atom.
Electronegativity increases from left to right across a period and decreases from top
to bottom down a group. Fluorine is the most electronegative atom. Francium is the
least electronegative.
Chapter 2 Notes for Chem Exam
Chapter 2: Atoms and Molecules
2.1 Symbols and Formulas:
• Element – homogenous pure substances made up of identical atoms
• 88 naturally occurring elements found in the Earth’s crust, oceans, and
atmosphere
• Each element can be characterized and identified by its unique set of
physical and chemical properties.
• Each element is therefore assigned a unique name and symbol, called an:
• Elemental Symbol – based on element’s name and consists of a single
capital letter or a capital letter followed by one lowercase letter.
Example: H – Hydrogen, He – Helium
• Compounds – pure substances made up of two or more different kinds of
atoms.
• Atoms are identical, whether in an element or compound, so the symbols
used for elements are the same used for the atoms in a compound.
• Compound Formula – Symbol for the molecule of a compound, consisting
of the symbols of the atoms in that compound.
Example: Hydrochloric Acid, (aka stomach acid), is written as HCl.
• Atoms present in numbers > 1 in a compound have that number indicated by
a subscript.
Example: Carbon Dioxide is written as CO2.
• A subscript number one is never used to indicate when only one atom is
present in a compound.
• The # 1 is implied if no other subscript is written in a compound formula.
Practice Problems: Write formulas for the following compounds.
1) Sulfuric Acid: two hydrogen atoms (H), one sulfur atom (S), and four oxygen
atoms (O)
2) Glucose: six carbon atoms (C), twelve hydrogen atoms, and six oxygen atoms.
2.2 Inside the atom:
• Atom – the limit of chemical subdivision of matter, the basic building block
of matter
• Atoms are made up of over 100 smaller (subatomic) particles
• Three major subatomic particles contribute the most influence on an atom’s
characteristics: Protons, Neutrons, and Electrons.
Table 2.3 Characteristics of important subatomic particles
Characteristics
Particle Symbol Charge Mass (g) Mass (u) Location
Electron e- -1 9.07 x 10-28 1/1836 Outside
Nucleus
Proton p, p+, H+ +1 1.67 x 10-24 1 Inside
Nucleus
Neutron n 0 1.67 x 10-24 1 Inside
Nucleus
• Nucleus – Central core of the atom. It is made up of neutrons and protons.
It contains 99.99% of the atom’s mass.
• Protons and neutrons are tightly bound together. Each nucleus has a positive
charge equal to the number of protons it contains.
• Even though the mass of a proton is 1836 times greater than the mass of an
electron, the charges of e- and p+ are of equal but opposite strength.
• So, an atom with equal #s of p+ and e- has no net charge and is considered to
be neutral.
• Electrons are negatively charged particles located outside of the nucleus.
http://www.silvershake.com/store/amethyst/images/Atomic-Structure.gif
• Electrons move very rapidly around the nucleus, throughout a relatively
large volume of space.
• Subatomic particles by themselves are relatively unstable, short-lived, and
do not display the properties of any element.
• The only way they gain long-term stability is by combining with other
particles to form an atom.
• Therefore atoms are considered the fundamental building blocks of matter.
Chem. 120/121 Chapter 2 Lecture Notes
2.3 Isotopes:
• Most atoms prefer to be neutral most of the time and are most stable when
they have no net charge (# of p+ = # of e-)
• Since neutrons have no charge, the number of neutrons can vary from the
number of protons and electrons in an atom.
• Atomic number (Z) – The number of protons in the nucleus of an atom.
• The atomic number is also the number of electrons in the neutral atom.
• ALL atoms of a specific element MUST have the same atomic number. (The
# of p+ in an atom is what give the element its identity)
• But, the number of neutrons can vary among atoms of the same element.
Example:
http://images.encarta.msn.com/xrefmedia/aencmed/targets/illus/ilt/T046738A.gif
Hydrogen (H) exists in 3 different atomic forms.
All 3 forms have the same atomic number (Z = 1 for 1 proton). They also all have
1 electron and all have zero net charge (They’re neutral).
Where they differ is these three forms contain 0, 1, and 2 neutrons respectively.
Note: all three have different names (protium, deuterium, and tritium) and different
properties (tritium is radioactive).
• Isotopes – atoms with the same atomic # but different numbers of neutrons.
• Mass number (A) – The sum total of the # of protons and # of neutrons in
the nucleus of the atom.
• Mass #s: Protium (A =1), Deuterium (A = 2), Tritium (A = 3).
• To distinguish between isotopes, the following notation is used:
• A
ZE where E is the elemental symbol, A is the mass #, and Z is atomic #
Example: 1
1H = protium, 2
1H = deuterium, 3
1H = tritium
Practice Problems:
1) What is the atomic #, mass #, and isotope symbol for an atom with 4 protons
and 5 neutrons?
2) How many neutrons are contained in an atom of chlorine-37?
2.4 Relative Masses of Atoms and Molecules:
• The masses of subatomic particles are very small and difficult to work with.
• Atomic Mass Unit (amu or u) – A unit used to express the relative masses
of atoms. One u is equal to 1/12th the mass of an atom of carbon–12.
• One atomic mass unit is ~ the weight of one proton or one neutron.
• Atomic weight – The mass of an average atom of an element expressed in
atomic mass units.
• Molecular weight (MW) – The relative mass of a molecule expressed in
atomic mass units and calculated by adding together the atomic weights of
the atoms in the molecule.
Ex: Water (H2O) has a MW of 18u: [2 x 1u (Z of H) + 16u (Z of O)] = 18u
Practice Problems: (1) Which element has atoms that are closest to twice the mass
of copper (Cu)?
(2) How many Helium (He) atoms would be required to have a mass ~ equal to the
mass of a single Neon (Ne) atom?
(3) What is the molecular weight of ethanol (C2H6O)?
2.5 Isotopes and Atomic Weights:
• Atomic Weight –the average mass of all of the atoms of a particular
element.
• Protons and Neutrons both have masses of 1u and the mass of e- are ~ 0.
• So, the 3 isotopes of H have different masses that are the sum of the p+ and
n in the nucleus of each atom (1 u, 2 u, 3 u).
• According to the Periodic Table, Hydrogen has an atomic weight of 1.008 u.
• Where does that # come from?
• It is the average weight of all H atoms.
Example: I have 3 different types of poker chips. I have 60 chips that weigh 11g,
30 chips that weigh 8g, and 10 chips that weigh 2g.
What is the average weight of my poker chips?
60% of my chips weigh 11g, 30% weigh 8g, and 10% weigh 2g, so multiply the
weight of each by its percentage and add the weights together.
11g (0.60) + 8g (0.30) + 2g (0.10) = 9.2g
Practice Problem:
1) Chlorine has two isotopes, 35Cl and 37Cl. 75.53% of all Chlorine is 35Cl (mass =
34.97 u) and 24.47% is 37Cl (mass = 36.97 u). Calculate the atomic weight.
2) Mg has 3 isotopes, 24Mg (23.99 u, 78.70%), 25Mg (24.99 u, 10.13%), and 26Mg
(25.98 u, 11.97%) Calculate the atomic weight.
Chem. 120/121 Chapter 2 Lecture Notes
2.6 – Avogadro’s Number - The Mole:
• Mole (mol) – The number of particles (atoms or molecules) contained in a
sample of element or compound with a mass in grams equal to the atomic or
molecular weight.
• 1 Mole = 6.022 x 1023 particles
• What the heck does that mean?
1 mol C atoms = 6.022 x 1023 C atoms = 12.01 g C
and
1 mol O atoms = 6.022 x 1023 O atoms = 16.00 g O
and
1 mol CO2=6.022 x 1023 CO2 molecules=44.01 g CO2
• How is that possible?
• Using modern scientific equipment, masses of indiviual atoms have been
determined.
• One atom of C has a mass of 1.99 x 1023 g
• One atom of O has a mass of 2.66 x 1023g
• One molecule of CO2’s mass is 7.31 x 1023g
Math check: 1.99g + 2.66g + 2.66g = 7.31g
How many atoms of C, O, and elements of CO2 would it take to equal the
Atomic or Molecular Masses?
• Using the Periodic Table, The atomic mass of C is 12.01u, the atomic mass
of O is 16.00u, and the molecular mass of CO2 is 44.01u.
12.01 g C 1 atom C = 6.02 x 1023 atoms C
1.99 x 10-23 g of C
16.00 g O 1 atom O = 6.02 x 1023 atoms O
2.66 x 10-23 g O
44.0g CO2 1 molecule CO2 = 6.02 x 1023 atoms CO2
7.31 x 10-23g CO2
• The number 6.022 x 1023 is a constant for all atoms or molecules.
• 6.022 x 1023 is ALWAYS the number of particles (atoms or molecules)
contained in a sample of element or compound with a mass in grams equal
to the atomic or molecular weight.
• A mole is a defined number, just like a dozen or a score. A dozen atoms =
12 atoms, a mole of atoms = 6.022 x 1023 atoms.
• Don’t get freaked out because it is a really big number and has a strange
name!
Practice Problems:
1) What is the mass in grams of 1.42 mol of Na?
2) How many moles of P atoms are in 67 g of P?
3) What is the mass in grams of one atom of N?
4) How many S atoms are in 98.6 g of S?
2.7 – The Mole and Chemical Formulas:
• The Chemical Formula of water is H2O.
• All water molecules contain H and O in a 2:1 ratio.
Using this ratio:
- 2 H2O molecules have 4 H atoms & 2 O atoms.
- 200 H2O molecules have 400 H atoms & 200 O atoms.
- 6.022 x 1023 H2O molecules have 12.04 x 1023 H atoms & 6.022 x 1023 O atoms.
- 1 mol of H2O molecules has 2 mol of H atoms & 1 mol of O atoms.
• The ratios of atoms in compounds will always equal the ratio of moles of
that atom in moles of that compound.
Example: 1 mol of H2SO4 has 2 mol H, 1 mol S, and 4 mol of O atoms.
However, 5 mol H2SO4 has 10 mol H, 5 mol S, and 20 mol O atoms.
• What kind of fun can we have with this?
Chem. 120/121 Chapter 2 Lecture Notes
Practice Problems:
1) How many mols of horns, hooves, and tails are there in 1 mol of cows? How
many in 1.5 mols?
2) How many mols of cave trolls could 3 mols of Legolas kill if it takes 5 arrows to
kill each troll and each Legolas has 100 arrows?
3) How many mol of each atom are found in 1 mol of CHCl3?
4) How many mol of each atom are found in 0.5 mol of glucose (C6H12O6)?
5) How many mol of each atom are found in 3 mol of baking soda (NaHCO3)?
2.1 Symbols and Formulas:
• Element – homogenous pure substances made up of identical atoms
• 88 naturally occurring elements found in the Earth’s crust, oceans, and
atmosphere
• Each element can be characterized and identified by its unique set of
physical and chemical properties.
• Each element is therefore assigned a unique name and symbol, called an:
• Elemental Symbol – based on element’s name and consists of a single
capital letter or a capital letter followed by one lowercase letter.
Example: H – Hydrogen, He – Helium
• Compounds – pure substances made up of two or more different kinds of
atoms.
• Atoms are identical, whether in an element or compound, so the symbols
used for elements are the same used for the atoms in a compound.
• Compound Formula – Symbol for the molecule of a compound, consisting
of the symbols of the atoms in that compound.
Example: Hydrochloric Acid, (aka stomach acid), is written as HCl.
• Atoms present in numbers > 1 in a compound have that number indicated by
a subscript.
Example: Carbon Dioxide is written as CO2.
• A subscript number one is never used to indicate when only one atom is
present in a compound.
• The # 1 is implied if no other subscript is written in a compound formula.
Practice Problems: Write formulas for the following compounds.
1) Sulfuric Acid: two hydrogen atoms (H), one sulfur atom (S), and four oxygen
atoms (O)
2) Glucose: six carbon atoms (C), twelve hydrogen atoms, and six oxygen atoms.
2.2 Inside the atom:
• Atom – the limit of chemical subdivision of matter, the basic building block
of matter
• Atoms are made up of over 100 smaller (subatomic) particles
• Three major subatomic particles contribute the most influence on an atom’s
characteristics: Protons, Neutrons, and Electrons.
Table 2.3 Characteristics of important subatomic particles
Characteristics
Particle Symbol Charge Mass (g) Mass (u) Location
Electron e- -1 9.07 x 10-28 1/1836 Outside
Nucleus
Proton p, p+, H+ +1 1.67 x 10-24 1 Inside
Nucleus
Neutron n 0 1.67 x 10-24 1 Inside
Nucleus
• Nucleus – Central core of the atom. It is made up of neutrons and protons.
It contains 99.99% of the atom’s mass.
• Protons and neutrons are tightly bound together. Each nucleus has a positive
charge equal to the number of protons it contains.
• Even though the mass of a proton is 1836 times greater than the mass of an
electron, the charges of e- and p+ are of equal but opposite strength.
• So, an atom with equal #s of p+ and e- has no net charge and is considered to
be neutral.
• Electrons are negatively charged particles located outside of the nucleus.
http://www.silvershake.com/store/amethyst/images/Atomic-Structure.gif
• Electrons move very rapidly around the nucleus, throughout a relatively
large volume of space.
• Subatomic particles by themselves are relatively unstable, short-lived, and
do not display the properties of any element.
• The only way they gain long-term stability is by combining with other
particles to form an atom.
• Therefore atoms are considered the fundamental building blocks of matter.
Chem. 120/121 Chapter 2 Lecture Notes
2.3 Isotopes:
• Most atoms prefer to be neutral most of the time and are most stable when
they have no net charge (# of p+ = # of e-)
• Since neutrons have no charge, the number of neutrons can vary from the
number of protons and electrons in an atom.
• Atomic number (Z) – The number of protons in the nucleus of an atom.
• The atomic number is also the number of electrons in the neutral atom.
• ALL atoms of a specific element MUST have the same atomic number. (The
# of p+ in an atom is what give the element its identity)
• But, the number of neutrons can vary among atoms of the same element.
Example:
http://images.encarta.msn.com/xrefmedia/aencmed/targets/illus/ilt/T046738A.gif
Hydrogen (H) exists in 3 different atomic forms.
All 3 forms have the same atomic number (Z = 1 for 1 proton). They also all have
1 electron and all have zero net charge (They’re neutral).
Where they differ is these three forms contain 0, 1, and 2 neutrons respectively.
Note: all three have different names (protium, deuterium, and tritium) and different
properties (tritium is radioactive).
• Isotopes – atoms with the same atomic # but different numbers of neutrons.
• Mass number (A) – The sum total of the # of protons and # of neutrons in
the nucleus of the atom.
• Mass #s: Protium (A =1), Deuterium (A = 2), Tritium (A = 3).
• To distinguish between isotopes, the following notation is used:
• A
ZE where E is the elemental symbol, A is the mass #, and Z is atomic #
Example: 1
1H = protium, 2
1H = deuterium, 3
1H = tritium
Practice Problems:
1) What is the atomic #, mass #, and isotope symbol for an atom with 4 protons
and 5 neutrons?
2) How many neutrons are contained in an atom of chlorine-37?
2.4 Relative Masses of Atoms and Molecules:
• The masses of subatomic particles are very small and difficult to work with.
• Atomic Mass Unit (amu or u) – A unit used to express the relative masses
of atoms. One u is equal to 1/12th the mass of an atom of carbon–12.
• One atomic mass unit is ~ the weight of one proton or one neutron.
• Atomic weight – The mass of an average atom of an element expressed in
atomic mass units.
• Molecular weight (MW) – The relative mass of a molecule expressed in
atomic mass units and calculated by adding together the atomic weights of
the atoms in the molecule.
Ex: Water (H2O) has a MW of 18u: [2 x 1u (Z of H) + 16u (Z of O)] = 18u
Practice Problems: (1) Which element has atoms that are closest to twice the mass
of copper (Cu)?
(2) How many Helium (He) atoms would be required to have a mass ~ equal to the
mass of a single Neon (Ne) atom?
(3) What is the molecular weight of ethanol (C2H6O)?
2.5 Isotopes and Atomic Weights:
• Atomic Weight –the average mass of all of the atoms of a particular
element.
• Protons and Neutrons both have masses of 1u and the mass of e- are ~ 0.
• So, the 3 isotopes of H have different masses that are the sum of the p+ and
n in the nucleus of each atom (1 u, 2 u, 3 u).
• According to the Periodic Table, Hydrogen has an atomic weight of 1.008 u.
• Where does that # come from?
• It is the average weight of all H atoms.
Example: I have 3 different types of poker chips. I have 60 chips that weigh 11g,
30 chips that weigh 8g, and 10 chips that weigh 2g.
What is the average weight of my poker chips?
60% of my chips weigh 11g, 30% weigh 8g, and 10% weigh 2g, so multiply the
weight of each by its percentage and add the weights together.
11g (0.60) + 8g (0.30) + 2g (0.10) = 9.2g
Practice Problem:
1) Chlorine has two isotopes, 35Cl and 37Cl. 75.53% of all Chlorine is 35Cl (mass =
34.97 u) and 24.47% is 37Cl (mass = 36.97 u). Calculate the atomic weight.
2) Mg has 3 isotopes, 24Mg (23.99 u, 78.70%), 25Mg (24.99 u, 10.13%), and 26Mg
(25.98 u, 11.97%) Calculate the atomic weight.
Chem. 120/121 Chapter 2 Lecture Notes
2.6 – Avogadro’s Number - The Mole:
• Mole (mol) – The number of particles (atoms or molecules) contained in a
sample of element or compound with a mass in grams equal to the atomic or
molecular weight.
• 1 Mole = 6.022 x 1023 particles
• What the heck does that mean?
1 mol C atoms = 6.022 x 1023 C atoms = 12.01 g C
and
1 mol O atoms = 6.022 x 1023 O atoms = 16.00 g O
and
1 mol CO2=6.022 x 1023 CO2 molecules=44.01 g CO2
• How is that possible?
• Using modern scientific equipment, masses of indiviual atoms have been
determined.
• One atom of C has a mass of 1.99 x 1023 g
• One atom of O has a mass of 2.66 x 1023g
• One molecule of CO2’s mass is 7.31 x 1023g
Math check: 1.99g + 2.66g + 2.66g = 7.31g
How many atoms of C, O, and elements of CO2 would it take to equal the
Atomic or Molecular Masses?
• Using the Periodic Table, The atomic mass of C is 12.01u, the atomic mass
of O is 16.00u, and the molecular mass of CO2 is 44.01u.
12.01 g C 1 atom C = 6.02 x 1023 atoms C
1.99 x 10-23 g of C
16.00 g O 1 atom O = 6.02 x 1023 atoms O
2.66 x 10-23 g O
44.0g CO2 1 molecule CO2 = 6.02 x 1023 atoms CO2
7.31 x 10-23g CO2
• The number 6.022 x 1023 is a constant for all atoms or molecules.
• 6.022 x 1023 is ALWAYS the number of particles (atoms or molecules)
contained in a sample of element or compound with a mass in grams equal
to the atomic or molecular weight.
• A mole is a defined number, just like a dozen or a score. A dozen atoms =
12 atoms, a mole of atoms = 6.022 x 1023 atoms.
• Don’t get freaked out because it is a really big number and has a strange
name!
Practice Problems:
1) What is the mass in grams of 1.42 mol of Na?
2) How many moles of P atoms are in 67 g of P?
3) What is the mass in grams of one atom of N?
4) How many S atoms are in 98.6 g of S?
2.7 – The Mole and Chemical Formulas:
• The Chemical Formula of water is H2O.
• All water molecules contain H and O in a 2:1 ratio.
Using this ratio:
- 2 H2O molecules have 4 H atoms & 2 O atoms.
- 200 H2O molecules have 400 H atoms & 200 O atoms.
- 6.022 x 1023 H2O molecules have 12.04 x 1023 H atoms & 6.022 x 1023 O atoms.
- 1 mol of H2O molecules has 2 mol of H atoms & 1 mol of O atoms.
• The ratios of atoms in compounds will always equal the ratio of moles of
that atom in moles of that compound.
Example: 1 mol of H2SO4 has 2 mol H, 1 mol S, and 4 mol of O atoms.
However, 5 mol H2SO4 has 10 mol H, 5 mol S, and 20 mol O atoms.
• What kind of fun can we have with this?
Chem. 120/121 Chapter 2 Lecture Notes
Practice Problems:
1) How many mols of horns, hooves, and tails are there in 1 mol of cows? How
many in 1.5 mols?
2) How many mols of cave trolls could 3 mols of Legolas kill if it takes 5 arrows to
kill each troll and each Legolas has 100 arrows?
3) How many mol of each atom are found in 1 mol of CHCl3?
4) How many mol of each atom are found in 0.5 mol of glucose (C6H12O6)?
5) How many mol of each atom are found in 3 mol of baking soda (NaHCO3)?
Chapter 1 Notes for Chemistry Exam
Chapter 1: Matter, Measurements, and Calculations
1.1 What is Matter?:
• Matter – anything that has mass and occupies space
• Mass – a measurement of the amount of matter in an object
• Weight – a measurement of the gravitational force acting on an
object
Mass is constant, weight varies depending on gravity
1.2 Properties of Matter and Changes in Matter:
The properties of matter are classified into two types:
• Physical – can be observed or measured without changing or
trying to change the composition of the matter in question
• Chemical – Properties matter demonstrates when attempts are
made to change it into new substances
Likewise, the changes matter undergoes is also classified into two
types:
• Physical Changes – Don’t change the composition of the
substance (reshaping, resizing)
• Chemical Changes – Convert matter into a new substance
(burning, digesting)
Dividing a substance into smaller and smaller amounts doesn’t change
its’ properties, so how small can you go?
1.3 A Model of Matter:
A molecule is the smallest part of a pure substance that has the
properties of that substance and is capable of a stable existence. It is the
limit of physical subdivision for a pure substance.
What if I do want to change the properties of the substance?
Can molecules be chopped into smaller pieces? - Yes, atoms
Atoms – The limit of chemical subdivision of matter, the basic building
block for all matter, the smallest particles of matter that can result from
chemical changes
John Dalton was the first to develop the atomic theory of matter. It
states:
— All matter is made up of tiny particles called atoms
— All atoms of a specific element are identical to each other and
different from atoms of a different element
— All compounds are combinations of atoms of two or more
elements
— Every molecule of a specific compound always contains the
same number of atoms of each kind of element found in the
compound
— In chemical reactions, atoms are rearranged, separated, or
combined, but are never created or destroyed
1.4 Classifying Matter:
Is it a pure substance or a mixture of substances?
Pure substances have constant compositions throughout, fixed
properties and can’t be separated into simpler substances
Mixtures are not constant throughout and can be separated into two or
more substances by physical means. Their properties vary with their
composition.
Pure substances and some mixtures (like salt water) can be homogenous
throughout.
Homogenous matter has a uniform appearance and the same properties
throughout.
A solution is a homogenous mixture of two or more substances.
Mixtures in which properties and appearance are not the same
throughout are called heterogeneous matter.
Element – A pure substance consisting of only one kind of atom.
Ex: gold (Au), oxygen (O2), graphite (C)
Compound – Pure substance consisting of two or more kinds of atoms.
Ex: water (H2O), carbon dioxide (CO2)
A compound can be chemically split into elements or other compounds.
1.6 The Metric System:
The official system of measurement used in science.
The metric system is a decimal system in which larger and smaller units
of quantity are all related to each other by factors of 10.
basic units of measurement – A specific unit from which other units
for the same quantity are obtained by multiplication or division.
As the amount increases or decreases, a prefix is added to the basic unit
to indicate a factor of 10 change from the basic unit.
Example: a 5K race is a distance of 5 kilometers, or 5 thousand meters.
The meter is the basic unit of distance in the metric system and the
prefix kilo means 1000, so 5 kilometers is 5,000 meters.
Common prefixes of the metric system (Table 1.2, pg 13):
mega (M) - 1 million times basic unit, 106 x basic unit
kilo (k) – 1 thousand times basic unit, 103 x basic unit
deci (d) – 1/10th the size of the basic unit, 10-1 x basic unit
centi (c) – 1/100th the size of the basic unit 10-2 x basic unit
milli (m) – 1/1000th the size of the basic unit 10-3 x basic unit
micro (μ) – 1/1,000,000 the size of the basic unit 10-6 x basic unit
nano (n) – 1/1,000,000,000 the size of the basic unit 10-9 x basic unit
pico (p) – 1/1,000,000,000,000 the size of the b-unit 10-12 x basic unit
Basic units for commonly used measurements: See table 1.3, pg 17
— Length – meter (m)
— Volume – cubic decimeter (dm3)
— Mass – kilogram (kg)
— Temperature – Kelvin (K)
— Energy – joule (J)
— Time – second (s)
All temperatures used in equations for this and all other science courses
must be converted into degrees Kelvin.
Depending on the scale, water boils at 212 oF, 100 oC, and 373 K; and
water freezes at 32 oF, 0 oC, and 273 K.
To convert between the different temperature scales:
oC = 5/9(oF-32) oF = 9/5(oC) + 32
oC = K – 273 K = oC + 273
Chem 120/121 Lecture Notes, Chapter 1. Friday, Jan. 16th
5
1.7 Large and small numbers:
When using really big or small numbers, scientific notation is used.
Scientific notation shows numbers as a product of a nonexponential term
and an exponential term, M x 10n.
The nonexponential term, M, is a whole number between 1 and <10,
written with the decimal to the right of the first non-zero digit. This is
the standard position.
The exponential term is a 10 raised to a whole number exponent that
may be either positive or negative. The value of n is the number of
places the decimal must be moved from the standard position in M to be
at the original position in the number when written normally.
If n is positive, the original position is to the right of the standard
position (the original number is bigger) and if n is negative, the original
position is to the left of the standard position (the original number is
smaller)
To multiply numbers in scientific notation, multiply the nonexponentials
and add the exponents.
Example: (2.2 x 103)(4 x 105) = (2.2 x 4)3+5 = 8.8 x 108
To divide numbers in scientific notation, divide the nonexponentials and
subtract the exponents.
Example: (9 x 107)/(2 x 105) = (9/2)7-5 = 4.52
1.8 Significant Digits:
All measurements contain a certain amount of uncertainty. The uncertainties are
determined by the limits of the measuring device.
Significant figures are the numbers in a measurement that represent the certainty of
the measurement, plus one more number representing an estimate.
Rules for determining the significance of zeros:
1. Zeros not preceded by nonzero numbers are not significant. 0.0037, 2 SDs
2. Zeros located between nonzero numbers are significant. 2.003, 4 SDs
3. Zeros located at the end of a number are significant. 4.0, 2 SDs
When using a measured number in a calculation, the answer can’t have more
certainty than the least certain measured value and should be written to reflect an
uncertainty equal to the most uncertain measurement.
Rules for multiplying and dividing:
The answer must contain the same # of SDs as the quantity with the fewest SDs.
Example: 10.456647447 x 2.0 = 21; Ex. 2) 150.101 / 10.05 = 15.03
Rules for Adding and subtracting: The answer obtained must have the same # of
digits to the right of the zero as the quantity with the fewest number of places to
the right of the zero.
Example: 8.73 + 4 = 13; Ex. 2) 55.120 – 10.11005 = 45.010
Calculators don’t use SDs, you must know how many SDs to use!
Not all numbers used in calculations have uncertainty, such as exact numbers,
these numbers don’t determine the number of SDs to be used in calculations.
For example, 1000 meters = 1 kilometer. This is by definition, there is no
uncertainty in this statement, therefore these numbers are not considered when
deciding how many SDs to use in an answer.
Another example of exact numbers are counting numbers, such as a dozen eggs =
12 eggs. No uncertainty, 1 dozen = 12 by definition, not 11.9 or 12.23. Therefore
1 or 12 in this example would not be used to determine the number of SDs to use.
1.9 How to Use Units in Chemistry Calculations:
Most calculations in chemistry will involve converting one numerical
value of some unit into another numerical value of another unit.
Example: how many grams are in 16.8 pounds? Or how many yards are
in 7 kilometers?
The key to solving these problems is the simple 4 step method shown
below:
Step 1) Write down the known or given quantity, including the
numerical value and the units.
Step 2) Leave some space and set the known quantity equal to the units
of the unknown quantity.
Step 3) Multiply the known quantity by one or more factors, such that
the units of the factors cancel each other out and generate the units of the
unknown quantity.
Step 4) AFTER the units are correct, do the math to calculate the
numerical value that goes with the units.
Example: how many yards are in 7 kilometers?
The known quantity is 7 km.
7 km 1000 m 1.094 yd = 7658 yds
1 km 1 m
What about SDs? We can only use 1 SD based on the 7 km, so the
answer is 8,000 yds, or 8 x 103 yds.
1.11 Density of Matter:
Even though two objects may be the same size, they may have
very different weights. Likewise, a pillowcase stuffed with
socks would weigh a lot more than the same pillowcase stuffed
with pennies. These differences in mass and volume result from
the unique densities of the objects.
Density = mass of an object divided by its volume, D = m/V
Densities of solids are often written as g/cm3, Densities of
liquids are often written as g/mL. Either way is correct because
1 mL = 1 cm3
Two objects of roughly the same size, such as a racquetball and
a billiards ball, have drastically different masses. The billiards
ball feels much heavier than the racquetball. Since both
volumes are ~ the same and the mass of the BB is > mass of the
RB, the density of the BB must be > the density of the RB.
Given the formula D = m/V and the values for any two of the
three properties, you should be able to solve for the unknown
quantity by rearranging the equation to solve for the unknown.
Example: The density of iron metal is 7.2 g/cm3, if you have a
900 mg sample of iron, what is the volume of that sample?
Step one, rearrange the equation D = m/V, to solve for V.
Multiplying both sides by V gives VD = m. Now divide both
sides by D to give: V = m/D
Using the equation V = m/D, we know m = 900 and D = 7.2,
so: V = 900/7.2 = 125 cm3
WRONG!!!! Use must write down the units!!!!
900 mg cm3 = 125 mg cm3
7.2 g g
mg cm3/g are NOT the units of Volume! You must FIRST get
your units right and the math will do itself!
We are solving for V, or Volume. The units of Volume are cm3.
We must get rid of the mg and g in the answer, but how?
What is the relationship between mg and g? 1 gram = 1000
milligrams
Can we use this relationship to eliminate the mg and gram units
in our calculation?
900 mg cm3 1 g = 0.125 cm3
7.2 g 1000 mg
Is this the correct number of significant figures?
No, the answer should have 2 SDs since the density is given as
7.2 g/cm3
Therefore, the correct answer for the volume of 900 mg of iron
powder is 0.13 cm3
1.1 What is Matter?:
• Matter – anything that has mass and occupies space
• Mass – a measurement of the amount of matter in an object
• Weight – a measurement of the gravitational force acting on an
object
Mass is constant, weight varies depending on gravity
1.2 Properties of Matter and Changes in Matter:
The properties of matter are classified into two types:
• Physical – can be observed or measured without changing or
trying to change the composition of the matter in question
• Chemical – Properties matter demonstrates when attempts are
made to change it into new substances
Likewise, the changes matter undergoes is also classified into two
types:
• Physical Changes – Don’t change the composition of the
substance (reshaping, resizing)
• Chemical Changes – Convert matter into a new substance
(burning, digesting)
Dividing a substance into smaller and smaller amounts doesn’t change
its’ properties, so how small can you go?
1.3 A Model of Matter:
A molecule is the smallest part of a pure substance that has the
properties of that substance and is capable of a stable existence. It is the
limit of physical subdivision for a pure substance.
What if I do want to change the properties of the substance?
Can molecules be chopped into smaller pieces? - Yes, atoms
Atoms – The limit of chemical subdivision of matter, the basic building
block for all matter, the smallest particles of matter that can result from
chemical changes
John Dalton was the first to develop the atomic theory of matter. It
states:
— All matter is made up of tiny particles called atoms
— All atoms of a specific element are identical to each other and
different from atoms of a different element
— All compounds are combinations of atoms of two or more
elements
— Every molecule of a specific compound always contains the
same number of atoms of each kind of element found in the
compound
— In chemical reactions, atoms are rearranged, separated, or
combined, but are never created or destroyed
1.4 Classifying Matter:
Is it a pure substance or a mixture of substances?
Pure substances have constant compositions throughout, fixed
properties and can’t be separated into simpler substances
Mixtures are not constant throughout and can be separated into two or
more substances by physical means. Their properties vary with their
composition.
Pure substances and some mixtures (like salt water) can be homogenous
throughout.
Homogenous matter has a uniform appearance and the same properties
throughout.
A solution is a homogenous mixture of two or more substances.
Mixtures in which properties and appearance are not the same
throughout are called heterogeneous matter.
Element – A pure substance consisting of only one kind of atom.
Ex: gold (Au), oxygen (O2), graphite (C)
Compound – Pure substance consisting of two or more kinds of atoms.
Ex: water (H2O), carbon dioxide (CO2)
A compound can be chemically split into elements or other compounds.
1.6 The Metric System:
The official system of measurement used in science.
The metric system is a decimal system in which larger and smaller units
of quantity are all related to each other by factors of 10.
basic units of measurement – A specific unit from which other units
for the same quantity are obtained by multiplication or division.
As the amount increases or decreases, a prefix is added to the basic unit
to indicate a factor of 10 change from the basic unit.
Example: a 5K race is a distance of 5 kilometers, or 5 thousand meters.
The meter is the basic unit of distance in the metric system and the
prefix kilo means 1000, so 5 kilometers is 5,000 meters.
Common prefixes of the metric system (Table 1.2, pg 13):
mega (M) - 1 million times basic unit, 106 x basic unit
kilo (k) – 1 thousand times basic unit, 103 x basic unit
deci (d) – 1/10th the size of the basic unit, 10-1 x basic unit
centi (c) – 1/100th the size of the basic unit 10-2 x basic unit
milli (m) – 1/1000th the size of the basic unit 10-3 x basic unit
micro (μ) – 1/1,000,000 the size of the basic unit 10-6 x basic unit
nano (n) – 1/1,000,000,000 the size of the basic unit 10-9 x basic unit
pico (p) – 1/1,000,000,000,000 the size of the b-unit 10-12 x basic unit
Basic units for commonly used measurements: See table 1.3, pg 17
— Length – meter (m)
— Volume – cubic decimeter (dm3)
— Mass – kilogram (kg)
— Temperature – Kelvin (K)
— Energy – joule (J)
— Time – second (s)
All temperatures used in equations for this and all other science courses
must be converted into degrees Kelvin.
Depending on the scale, water boils at 212 oF, 100 oC, and 373 K; and
water freezes at 32 oF, 0 oC, and 273 K.
To convert between the different temperature scales:
oC = 5/9(oF-32) oF = 9/5(oC) + 32
oC = K – 273 K = oC + 273
Chem 120/121 Lecture Notes, Chapter 1. Friday, Jan. 16th
5
1.7 Large and small numbers:
When using really big or small numbers, scientific notation is used.
Scientific notation shows numbers as a product of a nonexponential term
and an exponential term, M x 10n.
The nonexponential term, M, is a whole number between 1 and <10,
written with the decimal to the right of the first non-zero digit. This is
the standard position.
The exponential term is a 10 raised to a whole number exponent that
may be either positive or negative. The value of n is the number of
places the decimal must be moved from the standard position in M to be
at the original position in the number when written normally.
If n is positive, the original position is to the right of the standard
position (the original number is bigger) and if n is negative, the original
position is to the left of the standard position (the original number is
smaller)
To multiply numbers in scientific notation, multiply the nonexponentials
and add the exponents.
Example: (2.2 x 103)(4 x 105) = (2.2 x 4)3+5 = 8.8 x 108
To divide numbers in scientific notation, divide the nonexponentials and
subtract the exponents.
Example: (9 x 107)/(2 x 105) = (9/2)7-5 = 4.52
1.8 Significant Digits:
All measurements contain a certain amount of uncertainty. The uncertainties are
determined by the limits of the measuring device.
Significant figures are the numbers in a measurement that represent the certainty of
the measurement, plus one more number representing an estimate.
Rules for determining the significance of zeros:
1. Zeros not preceded by nonzero numbers are not significant. 0.0037, 2 SDs
2. Zeros located between nonzero numbers are significant. 2.003, 4 SDs
3. Zeros located at the end of a number are significant. 4.0, 2 SDs
When using a measured number in a calculation, the answer can’t have more
certainty than the least certain measured value and should be written to reflect an
uncertainty equal to the most uncertain measurement.
Rules for multiplying and dividing:
The answer must contain the same # of SDs as the quantity with the fewest SDs.
Example: 10.456647447 x 2.0 = 21; Ex. 2) 150.101 / 10.05 = 15.03
Rules for Adding and subtracting: The answer obtained must have the same # of
digits to the right of the zero as the quantity with the fewest number of places to
the right of the zero.
Example: 8.73 + 4 = 13; Ex. 2) 55.120 – 10.11005 = 45.010
Calculators don’t use SDs, you must know how many SDs to use!
Not all numbers used in calculations have uncertainty, such as exact numbers,
these numbers don’t determine the number of SDs to be used in calculations.
For example, 1000 meters = 1 kilometer. This is by definition, there is no
uncertainty in this statement, therefore these numbers are not considered when
deciding how many SDs to use in an answer.
Another example of exact numbers are counting numbers, such as a dozen eggs =
12 eggs. No uncertainty, 1 dozen = 12 by definition, not 11.9 or 12.23. Therefore
1 or 12 in this example would not be used to determine the number of SDs to use.
1.9 How to Use Units in Chemistry Calculations:
Most calculations in chemistry will involve converting one numerical
value of some unit into another numerical value of another unit.
Example: how many grams are in 16.8 pounds? Or how many yards are
in 7 kilometers?
The key to solving these problems is the simple 4 step method shown
below:
Step 1) Write down the known or given quantity, including the
numerical value and the units.
Step 2) Leave some space and set the known quantity equal to the units
of the unknown quantity.
Step 3) Multiply the known quantity by one or more factors, such that
the units of the factors cancel each other out and generate the units of the
unknown quantity.
Step 4) AFTER the units are correct, do the math to calculate the
numerical value that goes with the units.
Example: how many yards are in 7 kilometers?
The known quantity is 7 km.
7 km 1000 m 1.094 yd = 7658 yds
1 km 1 m
What about SDs? We can only use 1 SD based on the 7 km, so the
answer is 8,000 yds, or 8 x 103 yds.
1.11 Density of Matter:
Even though two objects may be the same size, they may have
very different weights. Likewise, a pillowcase stuffed with
socks would weigh a lot more than the same pillowcase stuffed
with pennies. These differences in mass and volume result from
the unique densities of the objects.
Density = mass of an object divided by its volume, D = m/V
Densities of solids are often written as g/cm3, Densities of
liquids are often written as g/mL. Either way is correct because
1 mL = 1 cm3
Two objects of roughly the same size, such as a racquetball and
a billiards ball, have drastically different masses. The billiards
ball feels much heavier than the racquetball. Since both
volumes are ~ the same and the mass of the BB is > mass of the
RB, the density of the BB must be > the density of the RB.
Given the formula D = m/V and the values for any two of the
three properties, you should be able to solve for the unknown
quantity by rearranging the equation to solve for the unknown.
Example: The density of iron metal is 7.2 g/cm3, if you have a
900 mg sample of iron, what is the volume of that sample?
Step one, rearrange the equation D = m/V, to solve for V.
Multiplying both sides by V gives VD = m. Now divide both
sides by D to give: V = m/D
Using the equation V = m/D, we know m = 900 and D = 7.2,
so: V = 900/7.2 = 125 cm3
WRONG!!!! Use must write down the units!!!!
900 mg cm3 = 125 mg cm3
7.2 g g
mg cm3/g are NOT the units of Volume! You must FIRST get
your units right and the math will do itself!
We are solving for V, or Volume. The units of Volume are cm3.
We must get rid of the mg and g in the answer, but how?
What is the relationship between mg and g? 1 gram = 1000
milligrams
Can we use this relationship to eliminate the mg and gram units
in our calculation?
900 mg cm3 1 g = 0.125 cm3
7.2 g 1000 mg
Is this the correct number of significant figures?
No, the answer should have 2 SDs since the density is given as
7.2 g/cm3
Therefore, the correct answer for the volume of 900 mg of iron
powder is 0.13 cm3
Monday, January 26, 2009
Microbiology 201C Notes
Lecture 011509
Change takes work.
Microbiology is the study of microbes.
__________________________________________
Lecture 012009
Science is a test of knowledge. “Have you asked Mother Nature?”
Francis Bacon was a writer who developed a new way of teaching and research when it came to science.
Created Deductive reasoning.
Law – well established; law of gravity
Theory – general issues to describe an idea that has been supported over time by the use of tests using the idea;
Hypothesis – just an idea, may or may not be correct
Warburg – won a nobel prize(international prize) Nobel invented dynamite; Warburg go the nobel prize for chemistry, a jew who lived in berlin during the reign of hitler. “All of my hypotheses were wrong”
Science was the major discovery of the 17th century. Science is progressive, new information comes up all the time.
Microbiology – dealing with small living things, as a science different from other sciences
Bacteria
Natural history – study of things such as trees/animals
Touching is a great way to transmit diseases. Contagion – touching to distribute
Miasma – microbes working which causes smell
Resistance to disease – contracting the disease but building an immunity to it
Dyes – colouring a pigment to change the colour
Optics – lenses used to increase the size of something viewed
Organic chemistry – carbon based chemistry
Dyes, optics, and organic chemistry came from the germans
Golden age of microbiology – from 1857-1900, best time at 1875, during that time, the etiology of microbiology was established at that time. The beginning combination of brains and technology to further research microbiology.
Etiology – study of causes of diseases
Klch – german, country doctor and vetenarian, anthrax, argemedy and petri dish; page 18
Reach Chapter 1
1. Organism present – if a disease occurs then the organism will be present.
2. Isolated in pure culture – not contaminated by other cultures
3. Inject/infect a normal animal – causing a disease by artificially infecting an animal
4. Re-isolate agent from animal – relocate the disease after dead.
____________________________________________________________________________
Lecture 012209
For test Tuesday::
Lectures, Assigned readings
Chapter 1, read everything up to page 26. It will be on the quiz.
Chapter 9, pages 184-186,
Not on this test – Chapter 17 pg 370-end of chapter, Diseases.
Formula for biology:: P= (N * V)/ R;
P – probability (of getting some disease);
N – number of organism to which you are exposed.
V – virulence( capacity/ability) pathos pathogens(pathogens is a path crafted to begin) – damage to physical;
R – resistance; immune system; ability of the host to fight again
(Delta N)/Time
Binary fission – organism simulates (osmotrophs) nutrients by moving across the cell membrane. Once it has reached a large enough size it divides.
Generation Time – the time it takes bacteria to double.
Page 185 –
Lag phase – division is slow, bacteria is getting used to situation and nutrients
Log/Exponential phase – division is expansion where the increases is greater due to accustomization of nutrients and environments.
Stationary phase – division is steady however, death is moving at the same rate due to nutrient limiting
Death phase – no division, death increasing due to lack of nutrient or some other reaction that is killing them.
Begin writing reflective essay on Chapter 11. “Controlling/Limiting Growth”
Salt prevents microbial growth, MREs – meals ready to eat are deadly.
Sugar prevents microbial growth, prevents water from growth
Change takes work.
Microbiology is the study of microbes.
__________________________________________
Lecture 012009
Science is a test of knowledge. “Have you asked Mother Nature?”
Francis Bacon was a writer who developed a new way of teaching and research when it came to science.
Created Deductive reasoning.
Law – well established; law of gravity
Theory – general issues to describe an idea that has been supported over time by the use of tests using the idea;
Hypothesis – just an idea, may or may not be correct
Warburg – won a nobel prize(international prize) Nobel invented dynamite; Warburg go the nobel prize for chemistry, a jew who lived in berlin during the reign of hitler. “All of my hypotheses were wrong”
Science was the major discovery of the 17th century. Science is progressive, new information comes up all the time.
Microbiology – dealing with small living things, as a science different from other sciences
Bacteria
Natural history – study of things such as trees/animals
Touching is a great way to transmit diseases. Contagion – touching to distribute
Miasma – microbes working which causes smell
Resistance to disease – contracting the disease but building an immunity to it
Dyes – colouring a pigment to change the colour
Optics – lenses used to increase the size of something viewed
Organic chemistry – carbon based chemistry
Dyes, optics, and organic chemistry came from the germans
Golden age of microbiology – from 1857-1900, best time at 1875, during that time, the etiology of microbiology was established at that time. The beginning combination of brains and technology to further research microbiology.
Etiology – study of causes of diseases
Klch – german, country doctor and vetenarian, anthrax, argemedy and petri dish; page 18
Reach Chapter 1
1. Organism present – if a disease occurs then the organism will be present.
2. Isolated in pure culture – not contaminated by other cultures
3. Inject/infect a normal animal – causing a disease by artificially infecting an animal
4. Re-isolate agent from animal – relocate the disease after dead.
____________________________________________________________________________
Lecture 012209
For test Tuesday::
Lectures, Assigned readings
Chapter 1, read everything up to page 26. It will be on the quiz.
Chapter 9, pages 184-186,
Not on this test – Chapter 17 pg 370-end of chapter, Diseases.
Formula for biology:: P= (N * V)/ R;
P – probability (of getting some disease);
N – number of organism to which you are exposed.
V – virulence( capacity/ability) pathos pathogens(pathogens is a path crafted to begin) – damage to physical;
R – resistance; immune system; ability of the host to fight again
(Delta N)/Time
Binary fission – organism simulates (osmotrophs) nutrients by moving across the cell membrane. Once it has reached a large enough size it divides.
Generation Time – the time it takes bacteria to double.
Page 185 –
Lag phase – division is slow, bacteria is getting used to situation and nutrients
Log/Exponential phase – division is expansion where the increases is greater due to accustomization of nutrients and environments.
Stationary phase – division is steady however, death is moving at the same rate due to nutrient limiting
Death phase – no division, death increasing due to lack of nutrient or some other reaction that is killing them.
Begin writing reflective essay on Chapter 11. “Controlling/Limiting Growth”
Salt prevents microbial growth, MREs – meals ready to eat are deadly.
Sugar prevents microbial growth, prevents water from growth
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