CHAPTER 2 – Building Blocks of Materials

Matter is everything that is around us that we interact with – “anything that occupies space and has mass”.

A classification scheme that you have likely seen….Chapter 2 1

Matter can generally be found in one of three forms:

Solids – rigid, occupying a defined volumeLiquids – fluid, occupying a defined volumeGases – fluid, occupying the space available

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We can explain the states of matter using “The Kinetic- Molecular Model of Matter” or “Kinetic-Molecular Theory”.

Indeed, it was the development of KMT that lead to the conclusion that atoms and molecules existed. The theory explains many different physical phenomena.

We will treat this in more detail in Chapter 11, but for now it is important to understand two fundamental concepts:

1)Matter exists in discrete units – either atoms, molecules, or ions.

2)Temperature is a measure of how fast atoms, molecules, or ions are moving.

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Homogeneous and Heterogeneous mixtures are classified based on whether or not we can “see” the distinct components. For example, salt in water is a homogeneous mixture as it is impossible to “see” the salt in the water.

Pepper in water is a heterogeneous mixture as the pepper floats and is easy to see. Indeed, pepper itself is a heterogeneous material made with many different types of material.

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Pure substances are composed of a single type of atom, molecule, or ionic compound.

We often speak of “pure water” or “pure gold” but these are actually only relative terms. It is virtually impossible to have a “pure” substance as there are almost always impurities present.

This has to do with the size of molecules. Consider that a single millilitre of water contains 3.34 x1022 molecules. To be pure, that would mean that all of these molecules must be water and that there is nothing else present.

Not likely. Instead, we speak of things being “pure” when we can no longer detect the impurities present.

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Levels of Operation – Levels of Understanding

“Observable level” – this is the level that we operate at in the laboratory and the world around us. We observe properties such as taste, smell, colour, and feel. Chemists add a whole layer of observation beyond these using spectroscopy.

“Molecular level” – this is the level that nature operates at in that everything is made up of atoms and molecules .

“Symbolic level” – this is the level that we use to connect the two other levels. We use symbolic representations to explain how molecular features lead to observable properties. We use symbolic representations when we discuss observable properties with other scientists.

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At both the symbolic and molecular levels, chemists employ “atoms” as the basic building block.

Literally, “atoms” means “not cuttable” . It is a term that originates in ancient Greece with a philosopher named “Demokritus of Abdera”.

Although you can not “see” atoms in the same sense that you can see the person next to you, we have many methods that allow us to “see” atoms and much evidence verifying their existence.

Hence, even though it is called “Atomic Theory” because, of course, it could be wrong, no scientist thinks in any other terms than atoms nor that there is another explanation.

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All atoms occur as one of the 114 known elements.

Realistically, there are only 90 or so naturally occurring elements and some of these are exceedingly rare.

The natural abundance of elements :

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Chapter 2 9

Note that atoms, molecules, and elements are not interchangeable terms.

An elements is composed of atoms of the same type. For example, elemental helium consists of atoms of helium and these are all single atoms.

Elemental oxygen, on the other hand, consists of oxygen molecules each of which has two oxygen atoms – it is O2 .

Note that we refer to both the element “oxygen” and “molecular oxygen” as “oxygen”. That is, air is predominantly nitrogen (N2 ) and oxygen (O2 ) but we would also say that carbon dioxide is “composed of carbon and oxygen” (CO2 ).

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Some elements occurring in their elemental form as atoms:

Helium (He) Neon (Ne)Argon (Ar) Krypton (Kr)Xenon (Xe) Radon (Rn)Mercury (Hg)

Some elements occurring in their elemental form as molecules:

Hydrogen (H2 ) Oxygen (O2 )Nitrogen (N2 ) Fluorine (F2 )Chlorine (Cl2 ) Bromine (Br2 )Iodine (I2 ) Phosphorus (P4 )Sulphur (S8 )

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All chemical compounds are composed of the 90 or so naturally occurring elements. The difference between one compound and another is:

a) which atoms – which elements – are present.b) the relative number of atoms of each element.c) how the atoms are connected to one another.

Compounds do not necessarily have the same chemical properties as their elements. NaCl is the example usually given:

Sodium is a shiny malleable metal which reduces water to give an explosive mixture of hydrogen.

Chlorine is a diatomic poisonous gas.Sodium chloride is a crystalline solid made of

sodium ions and chloride ions and found as table salt.

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+

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+

Chapter 2 14

Many compounds are “stoichiometric”(“stoichiometry” is the measure of a compound – it tells us the relationship between atoms/elements in a compound or compounds in a chemical reaction).

This leads to the concept of a chemical formula which shows the ratio of elements in a compound.

For example, carbon monoxide has a ratio of one carbon to one oxygen (CO), whereas carbon dioxide has one carbon to two oxygen atoms (CO2 ) . Carbonate has three oxygens for each carbon (CO3

2-) while percarbonate has four oxygens per carbon (CO4

2-). Each is a separate chemical species with its own observable properties and chemical characteristics.

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Chemical compounds can undergo transformation from one compound to another.

Consider carbon monoxide. It can pick up an oxygen atom to become carbon dioxide:

CO + ½ O2 CO2

The chemical reaction involves a change in the chemical compounds from reactants (those on the left hand side) to products (those on the right hand side).

During a chemical reaction, there is a redistribution of the atoms so that new substances are formed. The products have different properties from the reactants – they are different chemical compounds.

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Atoms are not actually the smallest component of matter. In the 1890s, it was realized that there are actually smaller “particles” which build atoms.

We now know that atoms are made up of electrons, protons, and neutrons.

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Atoms are composed of a nucleus containing protons and neutrons, surrounded by electrons. Protons have a positive charge and electrons have a negative charge – leading to electrostatic attraction between the two particles. Neutrons do not have a charge or are neutral.

Neutral atoms have equal numbers of protons and electrons. If an atom loses electrons, it can become a positively charged “cation”. If an atom gains electrons, it can become a negatively charged “anion”. The number of nuclear components (protons + neutrons) does not change in chemical reactions – just electrons.

The number of protons in the nucleus determines the value of the atomic number, Z, and is distinct for each element.

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Note: the electrons around the electrons do NOT occur in circles – they do not orbit around the nucleus.

The structure of the atom is much more complicated but much more explicable and the subject of quantum mechanics which we will get to in Section 8.4.

This model of the atoms dates back to Bohr (circa 1910) and is not current.

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The number of protons – the atomic number – defines which element an atom is. All carbon atoms, for example, have 6 protons. All oxygen atoms have 8 protons. All hydrogen atoms have only one proton.

But not all carbon, oxygen, or hydrogen atoms are the same. The number of neutrons that are found in the nucleus can vary. These atoms with the same atomic number but different mass are called “isotopes”.

For example, carbon-12 has 6 protons and 6 neutronscarbon-13 has 6 protons and 7 neutronscarbon-14 has 6 protons and 8 neutrons

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The mass number of an isotope is the sum of the atomic number and the number of neutrons:

Mass number, A = Z + number of neutrons= protons + neutrons

The atomic mass or “amu” reflects the mass number of the isotopes and the relative abundances of the different isotopes. For example,

98.89% of carbon atoms are 12C,1.11% are 13C,0.01% are 14C

as a consequence, Carbon is assigned a mass of 12.0107(8) amu – based on 12C having an exact mass of 12.

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Masses and Abundances of Some Isotopes, Atomic Weights of the Elements

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Atomic Weight (or Mass) is a weighted distribution based on the relative abundances:

At. Wt. = (% abundance of isotope 1) x (mass of isotope 1)+ (% abundance of isotope 2) x (mass of isotope 2)+ (% abundance…….) etcetera

For Hydrogen, this gives:

At. Wt. = 0.99985 x 1.0078 + 0.00015 x 2.0141 + 0 x 3.0161= 1.00794 amu

For Oxygen, this gives:

At. Wt. = 0.99757 x 15.99492 + 0.00038 x 16.99913 + 0.00205 x 17.99916 = 15.9994 amu

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How do we know what the weight of the isotopes is? How do we know their abundance?

Answer: Mass Spectroscopy

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Although everything around us is made up atoms, they are too small to be useful at the laboratory level. A mass spectrometer can count individual atoms, but it is impossible to do so on the bench top.

Consider that a single hydrogen atom would weigh 1.6735 x10-28 grams, it would be impossible to weigh out individual atoms for a chemical reaction.

Instead, we use a convenient quantity – the mole. This is the SI Unit for the amount of substance, which we will frequently designate by N.

“A mole of atoms of an element is the amount of it whose mass in grams is numerically equal to its atomic weight.”

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That is, “one mole of hydrogen atoms” is equal to 1.0107 grams. Note that “one mole of hydrogen” is equal to 2.0214 grams since we would normally find hydrogen as a diatomic molecule (H2 ).

The most important thing to realize is that one mole of hydrogen has the same number of particles as one mole of oxygen which has the same number of particles as one mole of water. That one mole is one mole.

That is, one mole is 6.0221415 x1023 somethings – be they molecules of water, atoms of iron, or chickens.

This is Avogadro’s constant, NA = 6.0221415 x1023 mol-1

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The concept of a mole allows us to work out the mass of molecules by simply adding up the atomic weights of the elements according to their appropriate stoichiometric factors.

That is, the mass of a molecule is simply the sum of the molar masses of its atoms.

Consider the reaction:

2H2 (g) + 1O2 (g) 2H2 O (g)

This means that “2 molecules of hydrogen react with 1 molecule of oxygen to give 2 molecules of water.”

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2H2 (g) + 1O2 (g) 2H2 O (g)

But it could also tell us that 2 moles of hydrogen gas react with 1 mole of oxygen gas to give us 2 moles of water.

Maybe, more importantly, it tells us that 4.0428 grams of hydrogen react with 31.9988 grams of oxygen to give 36.0416 grams of water.

All of these same ways of thinking say the same thing. They tell us the stoichiometric relationships for the reaction of hydrogen with oxygen to generate water.

We are operating at the symbolic level but it gives us something that is observable.

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In Figure 2.12, we have a comparison of the molar amount of a variety of substances:

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Information about atoms, atomic weights or mass, reactivity, and a whole bunch more is all contained in the Periodic Table.

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Chapter 2 32

Developing the Periodic Table

Many people contributed to the development of the Periodic Table – scientists that discovered the elements, scientists that tried to make sense of chemical properties, scientists that tried to organize chemical properties into an ordered picture.

However, in 1869, Dmitri Mendeleev published his thoughts in the form a table that recognized the periodic nature of the chemical properties known at the time.

He also recognized that there were elements that had not been discovered and left blanks in his table where these elements would occur while predicting their properties. For this reason, the Periodic Table is his invention.

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Mendeleev based his table on “atomic weights”. As the values for the weights become more refined, it became clear that there were some problems. Some elements were out of order (i.e. Ar and K).

In 1913, H.G.J. Moseley corrected Mendeleev’s assumption. His work on core x-rays led to precise measurement of the charge on the nuclei of the elements – and the supremacy of the atomic number (Z) as the organizing principle behind the periodic table.

The Law of Chemical Periodicity became:

The properties of the elements vary periodically with their atomic numbers.