Download Chapters 1-4 Numbers and Measurements in Chemistry Units SI

Chapters 1-4
What is Chemistry?
Chemistry is the study of the composition, structure, properties and
reactions of matter (the physical material of the universe). A main
challenge of chemistry is to bridge the macroscopic world we
experience with the microscopic structure of matter (atoms and
molecules).
Numbers and Measurements in Chemistry
• Chemists quantify data, expressing collected data with
units and significant figures.
– Units - designate the type of quantity measured.
– Prefixes - provide scale to a base unit.
– Significant Figures - indicate the amount of information that is
reliable when discussing a measurement.
Units
SI Prefixes
• The base unit designates
the type of quantity being
measured.
• SI units ((from French
Système International)
are the base units of
science.
• Some units comprise
combinations of these
base units and are
termed derived units
– 1 J = 1 kg m2 s-2
Dimensional Analysis
It may sound obvious (and yet is often forgotten)…many
mistakes can be avoided if an answer is checked for:
• units
• order of magnitude
• sign
• Prefixes are based on multiples of 10.
Dimensional Analysis Example
What is the mass of a cube of osmium that is 32 mm
on each side?
Conversion of a result from one system of units to another is
called dimensional analysis:
1. To carry out dimensional analysis, we must know the
relationship between units (equivalents): e.g. 1 nm = 10-9 m;
2. Use equivalents to determine unit factors: e.g. 1 = 1 nm/10-9 m;
3. Multiply result by appropriate unit factor(s) to convert units.
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Atomic Theory
• Each element is made up of tiny particles called
atoms.
• The atoms of a given element are identical; the
atoms of different elements are different in some
fundamental way or ways.
• Chemical compounds are formed when atoms
combine with each other. A given compound always
has the same relative numbers and types of atoms.
Atomic Theory
Atomic theory is the foundation of modern
chemistry….supported today by direct
visualization (and manipulation) of individual
atoms using scanning tunnelling microscopes
(STM):
(STM)
IBM spelled out in
individual atoms on
a surface…the first
“nano-advertisement”
• Chemical reactions involve reorganization of the
atoms - changes in the way they are bound together.
The atoms themselves are not changed in a chemical
reaction.
Fundamental Chemical Laws
Atomic Structure and Mass
• Law of conservation of mass
– Mass is neither created nor destroyed
• Law of definite proportion
– A given compound always contains exactly the
same proportion of elements by mass
• Law of multiple proportions
– When two elements form a series of compounds,
the ratios of the masses of the second element
that combine with 1 gram of the first element can
always be reduced to small whole numbers.
The Modern View of Atomic Structure
• The atom contains:
– electrons
– protons: found in the
nucleus; positive
charge equal in
magnitude to the
electron’s negative
charge.
– neutrons: found in the
nucleus; no charge;
virtually same mass as
a proton.
• The atom contains:
 Electrons – found outside the nucleus; negatively
charged.
 Protons – found in the nucleus; positive charge equal
in magnitude
ag ude to
o the
ee
electron’s
ec o s negative
ega e ccharge.
a ge
 Neutrons – found in the nucleus; no charge; virtually
same mass as a proton.
• The nucleus is:
 Small compared with the overall size of the atom.
 Extremely dense; accounts for almost all of the atom’s
mass.
Atomic Number and Mass Number
• Atomic Number, Z, is the number of protons in a nucleus.
– identifies the element
• Mass Number, A, is the sum of the number of protons and
number of neutrons in a nucleus.
• 1 amu = 1.6605 x 10-24 g
• Proton charge = 1.602x10-19 C
• Electron charge = -1.602x10-19 C
• Protons and neutrons are nearly 2000 times more massive than
electrons
Particle
Mass (amu)
Charge
Proton
1.007
+
Neutron
1.009
0
Electron
0.00055
–
• Heaviest atom is ~260 amu (4x10-22 g)
• Largest atom is 500 pm across.
• Typical C-C bond length 154 pm
2
Isotopes
• Isotopes are atoms of an element that differ in the
number of neutrons in their nucleus.
– same Z but different A
• Isotopic abundance is the mass percentage of an
isotope in a naturally occurring element.
• Show almost identical chemical properties; chemistry of
atom is due to its electrons.
• In nature most elements contain mixtures of isotopes.
Atomic Masses
• Entry for carbon on
the periodic table.
– Z=6
12
6
C
Atomic Symbols
• Information regarding atomic structure is
written in scientific shorthand called the atomic
symbol.
A
Z
E
– E is the atomic symbol for element
– Superscript A is the mass number.
– Subscript Z is the atomic number, it is redundant
and is often left off.
Atomic Masses
• Relative atomic mass for an element is an
average of the atomic masses for the
naturally occurring isotopes for an
element.
– Relative atomic mass
= 12.011 (~99%
carbon-12)
– Carbon-12 = 12.0000 x 0.9893 = 11.871 amu
– Carbon-13 = 13.0036 x 0.0107 = 0.1391 amu
– Element Symbol: C
– Average mass = 11.87 + 0.139 = 12.01 amu
Isotopes
Two Isotopes of Sodium
• Mass spectrum showing carbon isotopes.
3
Ions
• Ions are formed when the number of protons
and electrons in an atom are not equal.
– Ions with more protons than electrons are called
cations.
• net positive charge
– Ions with more electrons that protons are called
anions.
• net negative charge
• A monatomic ion is derived from a single
atom.
Ca2+, Cl-
• A polyatomic ion is derived from a group of
atoms with an overall charge.
NH4+, SO42-
Chemical Formulas
Chemical Formulas
• Chemical formulas describe a compound in
terms of the elements the compound
contains.
– The number of atoms for each element is
indicated by a subscript to the right of the
chemical symbol
symbol.
• Groups of atoms can be designated using parentheses.
Subscripts outside these parentheses mean that all
atoms enclosed in the parentheses are multiplied by the
value indicated by the subscript.
• Water molecules associated with certain compounds,
called hydrates, are indicated separately from the rest of
the compound.
Fe 3 (PO 4 )2  8H 2O
Chemical Formulas
• Compounds have different properties than their
constituent atoms.
• The molecular formula
for ethylene is C2H4.
• Ionic compounds contain cations and anions, usually
arranged in a lattice.
• The empirical formula
for ethylene is CH2.
• M
Molecular
l
l fformulas
l indicate
i di
the
h elements
l
and
d
number of atoms of each element actually contained
in a discrete unit of a compound.
• P
Polyethylene
l h l
can b
be
written as –[CH2CH2]n–
• Empirical formulas indicate the smallest whole
number ratio between the number of atoms of each
element in a molecular formula.
Chemical Bonding
– n is used to emphasize
that a large number of
these units are found in
an individual molecule
The Periodic Table
• All bonds are created by the exchange or sharing of
electrons.
• The Periodic Table is based on periodic law.
• The exchange or sharing of electrons results in lower
energy
gy for the compound
p
relative to the separate
p
atoms.
• Periodic law - when arranged properly, the
elements display a regular and periodic variation
in their chemical properties.
– Ionic: exchange; cations and anions present
– Metallic: sharing by forming a mobile “sea of electrons”
– Covalent: sharing between atoms
– Periods are horizontal rows on the periodic table.
– Groups are vertical columns on the periodic table.
4
Periodic Table
Periods and Groups
• Common names of specific groups:
–
–
–
–
Group 1: alkali metals.
Group 2: alkaline earth metals.
Group 17: halogens.
Group 18: noble gases/rare gases.
• Table Regions:
– Groups 1-2 and 13-18 are main group elements
– Groups 3-12 are transition metals
– Lanthanides and actinides are below the rest of
the table
Special Names for Grou
ups in the Periodic Table
Metals, Nonmetals, and Metalloids
• Metals are generally toward the left and bottom of the
periodic table. They are shiny, malleable, and ductile.
They conduct current and easily form cations.
• Nonmetals occupy the upper right-hand portion of the
periodic table. They are not shiny, malleable, or
ductile. They do not conduct current but do easily
form anions.
• Metalloids, or semimetals, have chemical properties
intermediate of metals and nonmetals. Metalloids are
clustered along a diagonal line on the periodic table
between the metals and nonmetals.
Organic Chemistry
Inorganic and Organic Chemistry
• Organic chemistry is the study of the compounds of
the element carbon, usually with oxygen, nitrogen,
and hydrogen.
– More than 18 million organic compounds exist.
– IIncludes
l d bi
biological
l i l molecules
l
l and
d nearly
l allll synthetic
th ti
polymers.
– Isomers: Different organic molecules that have the same
formula but are connected differently.
• Inorganic chemistry is the study of all other elements
and their compounds.
• Because carbon compounds can become quite large,
organic compounds are described simply and
unambiguously using line structures, where carbons
and hydrogens are not explicitly shown.
– Each corner or end of a line is a carbon.
– Hydrogen
y g atoms on carbon atoms are implied.
p
Carbon makes
four bonds, “missing” bonds go to hydrogen atoms. Hydrogen
can only make one covalent bond to another atom.
– Hydrogen atoms on any other element are shown
– All other elements are shown
CH 2
H 2N
OH
C
O
OH
becomes
H 2N
O
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Functional Groups
• Functional groups are arrangements of atoms that
tend to display similar chemical properties.
Alkanes
Alkanes have the general formula CnH2n+2 where n is an
integer.
– Chemical formulas are often written to emphasize functional
groups.
•
Methanol an alcohol
Methanol,
alcohol, is often written CH3OH instead of CH4O.
O
• Hydrocarbons contain only H and C atoms.
– Addition of functional groups to hydrocarbons results in
more complex compounds.
– Alkanes are hydrocarbons where the carbon atoms are
linked together with single bonds.
Functional Groups
Isomers
• Isomers are compounds
that have the same
chemical formula but are
connected differently.
– Three isomers of pentane,
C5H12.
• One straight chain
• Two branched chains
Functional Groups
Polymers
• Polymers are very large molecules made up of
many smaller molecules linked together.
• Monomers - The small molecules linked
together in polymers.
• Polymer backbone - The long chain of bonded
atoms formed when monomers link together to
form polymers.
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Polyethylene
Polyethylene
• The ethylene free radical reacts with another
ethylene monomer, extending the polyethylene chain.
• Ethylene monomers are linked together via a free
radical mechanism, which converts the carboncarbon double bond to a single bond.
– Free radicals have an unshared single electron and are
extremely reactive.
– An initiator produces the free radical that reacts with
ethylene, opening the double bond and transferring the
free radical to the ethylene monomer.
• Polymerization continues until the free radical reacts
with another free radical, or terminator, which
terminates the growth of the chain.
Polyethylene
Polymers
• Polyethylene
polymers can be
linear chains (highdensity polyethylene,
HDPE) or branched
chains (low-density
(low density
polyethylene, LDPE).
• Ultra-high molecular
weight polyethylene,
UHMWPE, contains
extremely long
chains.
• Polymers are the materials of choice for a host of
everyday objects.
a) polyethylene
b) polystyrene
c) poly(vinyl chloride), PVC
Chemical Equations
Polymers
•
• Models showing
how atoms are
arranged in
several
polymers.
•
•
A representation of a chemical reaction:
C2H5OH + 3O2  2CO2 + 3H2O
“reactants”
“products”
Reactants are only placed on the left side of the arrow,
products are only placed on the right side of the arrow.
The equation is balanced.
–
• Each of these
polymers has
distinct
properties.
•
•
•
All atoms present in the reactants are accounted for in the
products.
1 mole of ethanol reacts with 3 moles of oxygen to
produce 2 moles of carbon dioxide and 3 moles of water.
The balanced equation represents an overall ratio of
reactants and products, not what actually “happens”
during a reaction.
Use the coefficients in the balanced equation to decide
the amount of each reactant that is used, and the amount
of each product that is formed.
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Avogadro’s Number and the Mole
• A mole is a means of counting the large
number of particles in samples.
– One mole is the number of atoms in exactly 12
grams of 12C (carbon-12).
– 1 mole contains Avogadro’s number (6.022 x 1023
particles/mole) of particles.
– The mass of 6
6.022
022 x 1023 atoms of any element is
the molar mass of that element.
• Balanced chemical reactions also provide
mole ratios between reactants and products.
Determining Molar Mass
• The molar mass of a compound is the sum of
the molar masses of all the atoms in a
compound.
10gH 
1.0
16 0 g 
16.0

 2 mol H 
 +  1 mol O 

1 mol H  
1 mol O 
= 18.0 g/mol H 2O
2H2(g) + O2(g)  2H2O(g)
2 moles H2 : 1 mole O2 : 2 moles H2O
Calculations Using Moles and Molar Mass
Calculations Using Moles and Molar Mass
• Molar mass allows conversion from mass to number of
moles, much like a unit conversion.
• Avogadro’s number functions much like a unit conversion
between moles to number of particles.
– 1 mol C7H5N3O6 = 6.022  1023 C7H5N3O6 molecules
– 1 mol C7H5N3O6 = 227.133 g C7H5N3O6, (TNT)
300.0 g C 7 H 5 N 3O 6 
1 mol C 7 H 5 N 3O 6
227.133 g C 7 H 5 N 3O 6
= 1.320 mol C 7 H 5 N 3O 6
Molar Example
Potassium-40 is one of the few naturally occurring
radioactive isotopes of elements of low atomic number.
Its percent natural abundance among K isotopes is
0.012%. How many 40K atoms do you ingest by drinking
p of whole milk containing
g 371 mg
g of K?
one cup
– How many molecules are in 1.320 moles of TNT,?
1.320 mol C 7 H 5 N 3O 6 
6.022  10 23 molecules C 7 H 5 N 3O 6
1 mol C 7 H 5 N 3O 6
= 7.949  10 23 molecules C 7 H 5 N 3O 6
Fundamentals of Stoichiometry
• Stoichiometry is a term used to describe
quantitative relationships in chemistry.
– “How much?” of a p
product is p
produced or reactant is
consumed.
– Balanced chemical equation needed.
– Conversion between mass or volume to number of
moles frequently needed.
8
Ratios from a Balanced Chemical Equation
Ratios from a Balanced Chemical Equation
• This flow diagram illustrates the various steps
involved in solving a typical reaction stoichiometry
problem.
• Mole ratios are obtained from the coefficients in the
balanced chemical reaction.
– No different than unit conversion
– Usually more than one conversion is necessary
– 1 mol CH4 : 2 mol O2 : 1 mol CO2 : 2 mol H2O
– Write all quantities with their complete units
• These ratios can be used in solving problems:
2 mol H 2 O
1 mol CH 4
or
2 mol O2
1 mol CH 4
Ratios from a Balanced Chemical Equation
• How many grams of water can be produced if sufficient
hydrogen reacts with 26.0 g of oxygen?
Limiting Reactants
• In many chemical reactions, one reactant is
often exhausted before the other reactants. This
reactant is the limiting reactant.
– Limiting
Li iti reactant
t t iis d
determined
t
i d using
i stoichiometry.
t i hi
t
– The limiting reactant limits the quantity of product
produced.
Limiting Reactants
2H 2 (g) + O2 (g) 
 2H 2 O(g)
• Reaction between 6 H2 and
6 O2 will produce 6 H2O.
Introductory Examples
1. How many copper atoms are there on the end surface of 18-gauge
wire?
– 6 H2 can produce 6 H2O.
– 6 O2 can produce 12 H2O.
– H2 is limiting reactant.
– 3 O2 left over.
2. Oxygen can be generated by heating a solid if the oxygen is
loosely bound in the solid. KClO3 is an example of such a solid.
a. Write a balanced chemical equation for the decomposition of
KClO3 into KCl and O2.
b. How many grams of O2 can be produced from 10. g of KClO3?
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