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.

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 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:

Density

Densities of Some Common SubstancesSubstance Density (g/mL)Ice (0 °C) 0.917

Water (4.0 °C) 1.000Gold 19.31

Helium (25 °C) 0.000164Dry Air (25 °C) 0.001185

Human Fat 0.94Cork 0.22 - 0.26

Earth 5,54

Table Sugar 1,59Balsa Wood 0.12

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)

• 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 and 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:

 Substance  S (J/g °C) C (J/°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

Some common specific heats and heat capacities:

• 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 properties3) 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.

 Isotope  No. of atoms  Isotope

 Mass (amu)

 Xa  30  Xa  54 Xb  60  Xb  56Xc  10 Xc  59

 For Every 100 atoms of X

Isotope Masses of X

• 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.