Physical and Chemical Properties
• All substances have properties that we can use to identify
them. For example we can identify a person by their face,
their voice, height, finger prints, DNA etc.. The more of
these properties that we can identify, the better we know
the person. In a similar way matter has properties - and
there are many of them. There are two basic types of
properties that we can associate with matter. These
properties are called Physical properties and Chemical
properties:
• Physical properties:Properties that do not change the
chemical nature of matter
• Chemical properties:Properties that do change the
chemical nature of matter
• Examples of physical properties are: color, smell,
freezing point, boiling point, melting point, infra-red
spectrum, attraction (paramagnetic) or repulsion
(diamagnetic) to magnets, opacity, viscosity and density.
There are many more examples. Note that measuring each
of these properties will not alter the basic nature of the
substance.
• Examples of chemical properties are: heat of
combustion, reactivity with water, pH, and electromotive
force.
• The more properties we can identify for a substance, the
better we know the nature of that substance. These
properties can then help us model the substance and thus
understand how this substance will behave under various
conditions.
Conservation of Mass in Chemical Reactions
• Democritus (460-370 BC) and somewhat later John
Dalton (1766-1844) were the first to consider matter at its
most microscopic form. They both came up with the
concept of the "atom" as being the smallest unit of matter
and thus being undivisible*. This observation has an
important and fundamental consequence: mass is neither
created nor destroyed during the course of a chemical
reaction. How do we come to this conclusion? We know
that chemical reactions take place at the atomic/molecular
level. That is molecules and atoms interact with one
another during a chemical reaction. If atoms are
indivisible then they cannot be destroyed during a
chemical reaction. If atoms cannot be destroyed then the
mass of reactants must equal the mass of the products in a
chemical reaction. e.g.,
• Reactants -------> Products
• Mass of Reactants = Mass of Products
• This can be visualized by considering the formation of
water from oxygen and hydrogen molecules:
• Note that the hydrogen and oxygen atoms simply
rearrange themselves but are not destroyed. Therefore
mass is conserved.
Iron + Oxygen -----> Rust
100 g + ?g ------> 143g
mass reactants = mass products
mass products = 143g = mass reactants
= 100 + mass of oxygen
mass oxygen = 43 g
Elements, Compounds and Mixtures
• All substances have mass and therefore must be
composed of atoms. These atoms and how they assemble
themselves in the substance determines their chemical
and physical properties. Substances can be classified
according to how these atoms are assembled and is
known as Classification of Matter: All matter falls into
one of three categories: elements, compounds or
mixtures. Furthermore, mixtures can be classified as
homogeneous or inhomogeneous. The scheme looks
something like the diagram next slide:
• This classification depends upon how we try and separate
matter into its basic components. This separation is called
the "process". There are two processes: a physical and a
chemical process.
• Physical process: a process using physical properties
• Chemical process: a process using chemical properties
• If we have a sample of matter and can find a physical
process such as evaporation, magnets, color etc. to
separate it then the sample is a "mixture". Furthermore if
the sample is a mixture of solids and liquids (e.g., sand
and water) etc. or two or more liquids that don't mix (e.g.,
oil and vinegar) then the mixture is "inhomogeneous".
Otherwise the sample is a "homogeneous" mixture.
• If there is no physical process that will separate the
sample then the sample is a "pure" substance. If a
chemical process such as combustion or oxidation breaks
the substance down to its constituent atoms then the
substance is a "compound"(e.g., salt, sugar, water).
Otherwise the substance is an "element" (e.g., copper
penny, aluminum foil). Compounds are made up of
molecules or salts. Elements are made up of single types
of atoms.
Density
• Density is a physical property of matter. Most commonly
density refers either to the mass per unit volume (mass
density) or the number of objects (e.g., atoms, molecules)
per unit volume (number density). We will focus out
attention on mass density. The mass density has the units
mass/volume. Since volume has the units length "cubed"
then the SI unit of mass density is kg/m3. More common
units of density are g/ml or g/l. Substances have different
densities. In fact the density of a substance can often be
used to help identify it. In the next slide there is a table of
densities of common materials:
Densities of Some Common Substances
Substance
Density (g/mL)
Ice (0 °C)
0.917
Water (4.0 °C)
1.000
Gold
19.31
Helium (25 °C)
0.000164
Dry Air (25 °C)
0.001185
Human Fat
0.94
Cork
0.22 - 0.26
Table Sugar
1,59
Balsa Wood
0.12
Earth
5,54
An important example is water. The above table states that
liquid water has a mass of 1 g in every ml. Thus 2 ml of water
has a mass of 2 g etc.. Table sugar is more dense than water
by about 60 percent. Density does not depend upon size. For
example the water in a swimming pool has the same density a
glass of that swimming pool water.
Calculations with density are straight forward and involve the
formula for density namely D=m/V, where D=density,
m=
mass and V = volume.
• Example 1: What is the volume of a nugget of gold that
has a mass of 3.45 g? The density of gold can be looked
upon as a conversion factor from mass to volume i.e.,
• Example 2: A light substance is found to weigh 23 g and
to have a volume of 0.192 liters. What is the substance?
• Based upon this result we would guess that this substance
might be balsa wood.
• Example 3: What is the mass of 1 liter of sugar?
Compounds
•
•
•
There are two basic types of compounds. They are
distinguished by by the manner in which the atoms bind
to one another in the compound. These two types are
called "molecular" compounds and "salts" (or
equivalently "ionic" compounds):
Molecular compounds:These compounds are made up of
molecules whose atoms bind to one another through
"covalent" bonds.
Salts:The atoms in salts are held together with "ionic"
bonds. Unlike molecules, salts always form solids in a
regular array called a "crystalline solid".
• A bond is the "glue" that holds atoms together. In
compounds this glue can either be covalent or ionic.
• Covalent bonds:The electrons are shared between
atoms. Therefore this sharing of electrons provides the
glue.
• Ionic bonds:Ionic bonds occur due to the mutual
attraction between atoms with positive and negative
charges i.e., ions.
Examples of Molecules
N-hexane (top)
Acetaldeyhde
(top)
Taxol (left)
An Example of a Salt
Sodium Chloride (NaCl)
Energy and Chemical Reactions
• When matter undergoes transformations that
change its chemical and physical properties then
that transformation was brought about by a
chemical reaction. On the other hand chemical
reactions can only take place if there is sufficient
energy to make the reaction proceed. Therefore
energy is a prerequisite for chemical reactions.
• Energy can come in many forms e.g., heat, work, light, kinetic,
potential, chemical etc.. Moreover, energy can itself transform
among these various forms. For example a ball at the edge of a
table has zero kinetic energy and positive potential energy. If
the ball drops it will have zero potential energy and positive
kinetic energy the instant it hits the floor. However the sum of
the potential and kinetic energy is the same throughout the
ball's dropping history. Therefore energy has neither been
created or destroyed but has transformed from potential to
kinetic energy.
• Molecular of chemical energy can mean several things:
Chemical bonds are a source of energy, the movement of
molecules in space is kinetic energy, the vibrations and
rotations of molecules is another source of chemical
energy. All of these forms of chemical energy contribute
in one way or another to chemical reactions.
• The units of chemical reactions are straightforward and is
given in the diagram below:
• There are many other units for energy including electron
volt (ev), erg, kjoule (kJ) etc.
Specific Heat and Heat Capacity
• Specific heat is another physical property of matter. All
matter has a temperature associated with it. The
temperature of matter is a direct measure of the motion of
the molecules: The greater the motion the higher the
temperature:
• Motion requires energy: The more energy matter has the
higher temperature it will also have. Typically this energy
is supplied by heat. Heat loss or gain by matter is
equivalent energy loss or gain.
• With the observation above understood we can now ask the
following question: by how much will the temperature of an
object increase or decrease by the gain or loss of heat
energy? The answer is given by the specific heat (S) of the
object. The specific heat of an object is defined in the
following way: Take an object of mass m, put in x amount of
heat and carefully note the temperature rise, then S is given
by;
• In this definition mass is usually in either grams or
kilograms and temperatture is either in kelvin or degres
Celcius. Note that the specific heat is "per unit mass". Thus,
the specific heat of a gallon of milk is equal to the specific
heat of a quart of milk. A related quantity is called the heat
capacity (C). of an object. The relation between S and C is
C = (mass of obect) x (specific heat of object).
• A table of some common specific heats and heat
capacities is given below:
Some common specific heats
and heat capacities:
C (J/°C)
Substance S (J/g °C) for 100 g
Air
1.01
101
Aluminum 0.902
90.2
Copper
0.385
38.5
Gold
0.129
12.9
Iron
0.450
45.0
Mercury
0.140
14.0
NaCl
0.864
86.4
Ice
2,03
203
Water
4,179
417,9
• Consider the specific heat of copper , 0.385 J/g °C. What
this means is that it takes 0.385 Joules of heat to raise 1
gram of copper 1 degree Celsius. Thus, if we take 1 gram
of copper at 25 °C and add 1 Joule of heat to it, we will
find that the temperature of the copper will have risen to
26 °C. We can then ask: How much heat will it take to
raise by 1 °C 2g of copper?. Clearly the answer is 0.385 J
for each gram or 2x0.385 J = 0.770 J. What about a
pound of copper? A simple way of dealing with different
masses of matter is to determine the heat capacity C as
defined above. Note that C depends upon the size of the
object as opposed to S that does not.
• Example 1: How much energy does it take to raise the
temperature of 50 g of copper by 10 °C?
• Example 2: If we add 30 J of heat to 10 g of aluminum,
by how much will its temperature increase?
• Thus, if the initial temperature of the aluminum was
20 °C then after the heat is added the temperature will be
28.3 °C.
Dalton’s Atomic Theory
• Democritus first suggested the existence of the atom but it
took almost two millennia before the atom was placed on
a solid foothold as a fundamental chemical object by John
Dalton (1766-1844). Although two centuries old, Dalton's
atomic theory remains valid in modern chemical thought
. Dalton's Atomic Theory
1) All matter is made of atoms. Atoms are indivisible and
indestructible.
2) All atoms of a given element are identical in mass and
properties
3) Compounds are formed by a combination of two or
more different kinds of atoms.
4) A chemical reaction is a rearrangement of atoms.
• Modern atomic theory is, of course, a little more involved
than Dalton's theory but the essence of Dalton's theory
remains valid. Today we know that atoms can be
destroyed via nuclear reactions but not by chemical
reactions. Also, there are different kinds of atoms
(differing by their masses) within an element that are
known as "isotopes", but isotopes of an element have the
same chemical properties.
• Many heretofore unexplained chemical phenomena were
quickly explained by Dalton with his theory. Dalton's
theory quickly became the theoretical foundation in
chemistry.
Composition of the Atom
• Atoms have a definite structure. This structure determines
the chemical and physical properties of matter. This
atomic structure was not fully understood until the
discovery of the neutron in 1932. The history of the
discovery of atomic structure is one of the most
interesting and profound stories in science. In 1910
Rutherford was the first to propose what is accepted
today as the basic structure of the atom. Today the
Rutherford model is called the "planetary" model of the
atom. In the planetary model of the atom there exists a
nucleus at the center made up of positively charged
particles called "protons" and electrically neutral atoms
called "neutrons". Surrounding or "orbiting" this nucleus
are the electrons. In elements the number of electrons
equals the number of protons.
• The picture above greatly exaggerates the size of the
nucleus relative to that of the atom. The nucleus is about
100,000 times smaller than the atom. Nevertheless, the
nucleus contains essentially all of the mass of the atom. In
order to discuss the mass of an atom we need to define a
new unit of mass appropriate to that of an atom. This new
unit of mass is called the "atomic mass unit" or amu. The
conversion between the amu and gram is
1 amu = 1.67x10-24 g
• The mass, in amu, of the three particles is given in the
table below:
• Note that the mass of an electron is about 2000 times
smaller than that of the proton and neutron. Also note that
the mass of the proton and neutron is close to 1 amu. This
is a useful fact to remember. If the number of electrons
does not equal the number of protons in the nucleus then
the atom is an ion:
• cation: number of electrons < number of protons
• anion: number of electrons > number of protons
Rutherford’s Planetary Model of the Atom
• By 1911 the components of the atom had been
discovered. The atom consisted of subatomic
particles called protons and electrons. However,
it was not clear how these protons and electrons
were arranged within the atom. J.J. Thomson
suggested the"plum pudding" model. In this
model the electrons and protons are uniformly
mixed throughout the atom:
• Rutherford tested Thomson's hypothesis by devising his
"gold foil" experiment. Rutherford reasoned that if
Thomson's model was correct then the mass of the atom
was spread out throughout the atom. Then, if he shot high
velocity alpha particles (helium nuclei) at an atom then
there would be very little to deflect the alpha particles. He
decided to test this with a thin film of gold atoms. As
expected, most alpha particles went right through the gold
foil but to his amazement a few alpha particles rebounded
almost directly backwards.
• These deflections were not consistent with
Thomson's model. Rutherford was forced to
discard the Plum Pudding model and reasoned
that the only way the alpha particles could be
deflected backwards was if most of the mass in
an atom was concentrated in a nucleus. He thus
developed the planetary model of the atom which
put all the protons in the nucleus and the
electrons orbited around the nucleus like planets
around the sun.
Isotopes and Atomic Symbols
• Atomic Symbols:
The atom of each element is made up of
electrons, protons and neutrons. All atoms of the
same neutral element have the same number of
protons and electrons but the number of neutrons
can differ. Atoms of the same element but
different neutrons are called isotopes. Because of
these isotopes it becomes necessary to develop a
notation to distinguish one isotope from another the atomic symbol.)
The atomic symbol has three parts to it:
• 1. The symbol X: the usual element symbol
• 2. The atomic number A: equal to the
number of protons (placed as a left subscript)
• 3. The mass number Z: equal to the number
of protons and neutrons in the isotope (placed
as a left superscript
• Examples 1:
• Consider two isotopes of gallium, one having the 37
neutrons and the other having 39 neutrons. Write the
atomic symbols for each isotope. Solution:
• Example 2:
• How many neutrons does the isotope of copper with mass
number Z = 65 have?
Solution: From the periodic table we see that copper has
an atomic number of 29. Since Z is the number of protons
plus the number of neutrons, then No. neutrons = 65 - 29
= 36
The Mass
• The standard for every unit must be defined. Length is an
example. The basic unit of length is the meter which was
defined in 1983 as equal to the distance traveled by light
in a vacuum in 1/299,792,458 of a second. Mass must
also be defined. The definition of mass today is the amu
(atomic mass unit). The amu is defined in the following
way: the mass of one atom of the carbon-12 isotope is
EXACTLY 12 amu.
mass of one carbon-12 atom = 12 amu
• All other masses are measured relative to this carbon-12
standard. For example, suppose we do an experiment and
find that the isotope bromine-81 has a mass that is 6.743
times that of carbon-12. Then the mass of bromine-81
would be given by
Atomic Weights
• Most elements can be found on earth (with the exception
of those elements that too unstable and thus must be
synthesized in the laboratory). Since all elements have
isotopes then we must consider how much of one isotope
of an element exists versus another isotope of the same
element. These are called the "natural" abundances on
earth.
• Natural Abundances:
• Suppose we go to a cave and mine element "X".
After careful analysis we find that in our sample
of element X there exists three isotopes: Xa, Xb
and Xc. Moreover, we find that out of every 100
atoms the various isotopes are distributed as
follows, and their masses are given.
For Every 100
atoms of X
No. of Isotope
atoms
Xa
30
Xb
60
Xc
10
Isotope Masses of
X
Mass Isotope
(amu)
Xa
54
Xb
56
Xc
59
• Then the average mass (atomic weight) is given by:
• The atomic weight of each element is included along with
the element symbol in the periodic table. It is important to
note that no one atom has a mass equal to that of the
atomic weight. Remember: the atomic weight represents
that average mass of the atoms.