The Periodic Table and the Elements

CHEMISTRY 30S – MODULE 3 CHEMICAL REACTIONS

LESSON 1  Isotopes, Ions, and the Periodic Table

The processes in our bodies that allow us to breathe, digest food and convert food to energy are all a result of chemical reactions. The burning of natural gas to keep our houses warm in the winter and the burning of gasoline in our cars are chemical reactions. Chemical reactions occur everywhere in the world around us. They play a large role in our lives. In this module we will examine several types of chemical reactions and the quantities chemists use to describe chemical reactions.

Hopefully you have learned about the basic structure of the atom and have used the mass number and the atomic number to determine the number of neutrons in an atom, using the information on the periodic table.

When you have completed this lesson, you will be able to:  Define average atomic mass with respect to isotopes and relative abundance.

 Research the importance and applications of isotopes

 Use the periodic table to determine the subatomic particles that make up atoms

 Explain why atoms bond to make compounds, and define ions

All atoms (except hydrogen) are made of 3 basic particles: protons, neutrons and electrons. Protons are Positively charged, electrons are negatively charged, and NEUTRons are NEUTRal 

The mass of atoms come from the nucleus, which contains all the protons and neutrons, while electrons circle around the outside of the atom. The electrons have no mass compared to the protons and neutrons, so we consider the total mass of an atom to be contained in the sum of its protons and neutrons.

Each element has its own number of protons, which is its atomic number (we use Z to represent atomic number).

For example, the number of protons in Lithium is 3, and its atomic number is 3. In Hydrogen, (atomic number 1) there is one proton.

Thus we can tell the number of protons in an element’s atoms, by looking at the periodic table, and noting the atomic number of the element. And since atoms all have the same number of electrons as they have of protons, that atomic number is also equal to the number of protons.

CHEM 30S Mod 3 Chemical Reactions Lesson 1 1 TALC 2010

The number of neutrons in each atom varies, even between atoms of the same element. For example, potassium can exist as three different kinds of atoms. All three kinds of potassium atoms contain 19 protons, but one kind has 20 neutrons, one kind has 21 neutrons and yet another has 22 neutrons. Atoms that have the same number of protons but differ in their number of neutrons are called isotopes. Most elements exist as more than one isotope.

If different isotopes have different numbers of neutrons, they will also have different masses. The atomic mass unit (often designated as u, μ, or amu) is defined as 1/12 the mass of a Carbon atom. Why Carbon? Because it is a very common element. The amu is also the mass of a proton and of a neutron.  Chemists have designed a symbol for each isotope that includes the element’s symbol, its atomic number (Z) and its mass number (A).

Different variations of atoms of the same element occur in nature. These variations are called isotopes. The average mass of the isotopes for each element is a characteristic of that element.

Isotopes are atoms of the same element (which means they have the same number of protons) with different numbers of neutrons. They have identical atomic numbers (number of protons) but different mass numbers (number of protons plus number of neutrons).

If we consider the potassium isotopes previously mentioned, the isotope containing 19 protons and 20 neutrons will have a mass number of 39 (19 + 20). We call this isotope potassium-39. The isotope that has 19 protons and 21 neutrons will have a mass number of 40 (19 + 21) and is called potassium-40.

Z = atomic number = number of protons

A = mass number = number of protons plus number of neutrons

Number of neutrons is found by A – Z (mass number – atomic number)

CHEM 30S 2 M2 L1 TALC 2010

Isotopes are usually represented in a few different ways:

Example:

The symbol for potassium-39 would be:

The symbol for potassium-40 would be:

The relative abundance of an isotope is the fraction of each isotope found in an average sample of the element. Look at the periodic table, and notice that the atomic mass shown for each element on a periodic table is rarely a whole number; there are decimals following the numbers. If each proton and neutron have an atomic mass of 1 amu, then how does this happen? This is because it is actually an average mass of all isotopes of that element. Complete the following chart, assuming the most common forms of the atoms, rather than isotopes: COMPOSITION OF SELECTED NEUTRAL ATOMS Symbol Ba Cr Ag Atomic # 36 Mass # 12 # of protons 10 19 # of neutrons 4 # of electrons

Calculating Average Atomic Mass

Example 1. Mass spectrometers have shown that magnesium exists as three isotopes: magnesium-24, magnesium-25 and magnesium-26. In an average sample of magnesium, 78.99% is magnesium-24, 10.00% is magnesium-25 and 11.01% is magnesium-26. If the atomic mass of magnesium-24 is 23.985 amu, magnesium-25 is 24.986 amu and magnesium-26 is 25.982 amu, calculate the the average atomic mass of magnesium. Solution You will notice that the actual atomic mass of magnesium-24 is not 24 amu, but 23.985 amu and similarly for the other isotopes. This is because a single amu is not exactly equal to the mass of a proton or a neutron. The average atomic mass is the weighted average of the relative abundances of each isotope. We can think of the weighted average in terms of how your marks are determined

CHEM 30S 3 M2 L1 TALC 2010 in this, or any course. Tests may be worth 30%, quizzes may be worth 15%, labs and assignments may be 25% and the final exam 30%. For magnesium, that would be:

If you look at the periodic table, you will see that the average atomic mass of magnesium is about 24.3 amu. You will also notice that the average mass of the magnesium is closer to 24 amu than 25 amu or 26 amu. This is because the magnesium-24 has a greater abundance than the other two isotopes.

Radioisotopes

The nuclei of some isotopes can be unstable. As a result, the nuclei may release energy and/or particles. These type of isotopes are called radioisotopes or radioactive isotopes. As the energy/particles are released from the nucleus, the atom may be converted into another isotope of the same element or an isotope of a completely different element. This release of energy is called radioactive decay or just decay. There are several types of decay related to the types of particles released. We will not discuss these types of decay because it is beyond the scope of this course. If you are interested, there are many internet websites and textbooks you can use to investigate this further. A term often associated with radioisotopes is half-life. Half-life refers to the amount of time it takes for half the radioisotope to be converted into another isotope. For example, 238U has a half-life of 4.46 x 109 years. This means for a 10-gram sample of 238U it takes 4.46 x 109 years for half of it, or 5 grams, to decay into another isotope. Another way to think of half life is after 4.46 x 109 years only 5 grams of the original 10-gram sample of 238U will remain. Not all isotopes have such a long half-life; chlorine-38 has a half-life of 87.3 minutes, while element 106 has a half-life of 0.8 seconds.

Exercise: 1. Complete the following table to calculate the average atomic mass of each element.

Relative Average Mass Element Symbol Mass (µ) Abundance Atomic Mass Number (%) (µ) Carbon C-12 12 12(exactly) 98.98 C-13 13 13.003 1.11 Silicon Si-28 28 27.977 92.21

Si-29 29 28.976 4.70 Si-30 30 29.974 3.09

2. Define the term isotope. Explain how an element’s atomic mass is related to the abundances of its different isotopes.

3. Using the graph below, calculate the average atomic mass of copper.

The Periodic Table as a Classification System

Jacob Berzelius suggested the chemical symbols that everybody uses today. Berzelius used letters to represent the atoms of each element. For example: C for carbon, H for hydrogen, I for iodine, O for oxygen, P for phosphoprus, S for sulphur, N for nitrogen and F for fluorine.

With over 100 elements and only 26 letters in the alphabet it became necessary to include a second letter with the first. The second letter of the symbol is usually the second letter in the name of the element or a main consonant in the name. The first letter of the symbol is always written as an upper case or capital letter. The second letter is always written as a lower case letter (i.e., Ca not CA). The full name of the element is written with lower case (noncapital) letters only. For example:

Al -- aluminum, Ba -- barium, Br -- bromine, Ca -- calcium, Mg -- magnesium,

Zn -- zinc, Cl -- chlorine and As -- arsenic.

The name and symbol for all new elements is now established by IUPAC (The International Union of Pure and Applied Chemistry). IUPAC generally respects the recommendation(s) of the scientist(s) who discovered the element. IUPAC now requires two letters for any new element symbol.

The individual particles making up compounds (and certain elements) are referred to as molecules. If you have a glass of water, a container of oxygen gas, or a cube of sugar, you have a substance which consists of molecules. A molecule can be defined as a collection of two or more atoms held together strongly enough to form an individual particle. The attractive forces holding the atoms together in a molecule are known as chemical bonds. There is no absolute limit to the size of a molecule. Some molecules, like hydrogen chloride (HCl) and water (H2O) consist, respectively, of only two and three atoms. Other molecules are so large that they are sometimes referred to as “giant molecules”. Proteins are examples of such giant molecules. Although most elements are composed of individual atoms, a few elements naturally consist of pairs of atoms. These elements are made up of diatomic molecules – two atoms per molecule. It is the diatomic molecule that has the properties of the element. For example,the element oxygen consists of diatomic molecules. The molecule oxygen is represented by O2 When we inhale oxygen, along with the other gases in air, we inhale O2 molecules and not O atoms. ln fact, individual O atoms do not exist permanently in the atmosphere. All the properties of oxygen are due to the O2 molecules.

Compounds are formed when atoms combine. The individual particles making up the compounds are molecules. Since molecules are composed of atoms, it follows that once the masses of the atoms are known it is possible to determine the mass of a molecule. This mass is known as the molecular mass. This will be covered in module 2.

In chemical reactions, a rearrangement of only the electrons takes place -- the nucleus is not affected. Understanding chemistry depends upon understanding how the electrons of atoms are arranged around their nuclei and upon how the electrons of atoms interact with the electrons of other atoms.

When atoms approach each other closely, their electrons become simultaneously attracted by the positive nuclei of other atoms. The simultaneous attraction for the same electrons

CHEM 30S 6 M2 L1 TALC 2010 by the nuclei of two or more atoms causes electron rearrangements among atoms. The electron rearrangements may be considered to be of two distinct types:

1. the loss and gain of electrons 2. the sharing of electrons

The electron rearrangement is normally called “bonding”, and it results in chemical changes, and the formation of different compounds. This will be discussed in more detail later.

Simple Ions

In chemical reactions, atoms may lose or gain electrons to acquire the more stable electron structure of the nearest noble gas. The atoms lose or gain electrons and acquire an electron structure with the maximum possible number of electrons (i.e., 2, 8 or 18 electrons) in their outermost energy level.

The loss or gain of electrons results in a more stable electron energy level structure. The loss or gain of electrons also unbalances the number of positive (proton) and negative (electron) charges. Atoms are neutral (zero net charge) because of an equal number of positive (proton) and negative (electron) charges. Ions are charged (have a net charge) because of a different number of protons than electrons. The difference in the number of protons and electrons results from a loss or gain of electrons; the number of protons will never change in a chemical reaction.

Make sure you have a periodic table in front of you for this section. All of the simple ions in the first three periods (rows of the periodic table) will have the same number of electrons as the nearest noble gas. A noble gas is one of the elements in the far right column (family) of the periodic table. Elements close to the noble gases will try to obtain the electron structure (have the same number of electrons, whether they need to gain or lose) of the nearest noble gas. Each occupied energy level will contain the maximum number of electrons (i.e., 2, 8 and 18)

The name of a nonmetallic ion ends in “ide” while the name for a metallic ion uses just the name of the metal (i.e., chloride ion and magnesium ion). Metallic vs. nonmetallic will be discussed further in lesson 3. For now, notice the “zig-zag line” in the periodic table, towards the right hand side of the table. The elements to the left of the line are metals, the elements to the right are nonmetals. Memorize this fact!

Hydrogen atoms may either gain or lose an electron to form a one positive ion or a one negative ion.

Simple ions of the transition elements cannot be explained by the Bohr Model of the atom. The quantum mechanical model of the atom explains the transition metal (the

metals in the “dip” of the table, from Sc to Zn and those under them) ions in terms of energy sublevels. This is covered in later chemistry courses and is beyond the scope of this course.

LESSON 1 ASSIGNMENT 2

1. The extranuclear region of the atom, which makes up most of the volume of the atom, is occupied by ______2. Nearly all of the mass of any atom is made up of______and ______3. An atom has 53 protons in its nucleus. In a neutral atom it will also have______electrons and it will (gain I lose) ______(number)______electron(s) to acquire the electron population of the nearest noble gas, ______4. The atomic number of a K atom is (greater / less) than the atomic number of a Na atom. 5. The name of the ion formed by a bromine atom is ______6. The name of the ion formed by a calcium atom is ______7. The number of______in the nucleus of chlorine atoms may vary. 8. Atoms with the same number of protons but with a different number of neutrons in the nucleus are called ______9. The average mass of atoms for a particular element is called the ______10.Be sure you are able to define the following terms:

a) element b) compound c) electron d) proton e) neutron f) nucleus g) atom h) valence electron i) simple ion j) group k) period l) atomic number m) atomic mass n) isotope

CHEM 30S 8 M2 L1