CHM 101 Textbook Notes Ch 2

CHEM 101 – General Chemistry I
Chapter 2 Reading Assignment
Pages 31 - 60

2.1 – What is Matter Made Of?
Who 1st discovered atoms – Democritus where the first person to discover atoms. His followers would often think that there were multiple kinds of atoms and not just plain atoms.

2.2 – How Do We Classify Matter?
Elements – A substance that consists of identical atoms. Today there are 116 known elements. Example – C= Carbon; H= Hydrogen
Compounds – A pure substance made up of two or more elements in a fixed ratio by mass. Example – Water is made up of hydrogen and oxygen, H20.
Mixtures – A combination of two or more pure substances. Example – Blood, butter, gasoline, soap, the metal in a ring, the air we breathe, and the Earth.

2.3 – What are the Postulates of Dalton’s Atomic Theory?
Five Postulates 1. All matter is made up of very tiny, indivisible particles, which Dalton called atoms. 2. All atoms of a given element have the same chemical properties. Conversely, atoms of different elements have different chemical properties. 3. In ordinary chemical reactions, no atom of any element disappears or is changed into an atom of another element. 4. Compounds are formed by the chemical combination of two or more different kinds of atoms. In a given compound, the relative numbers of atoms of each kind of element are constant and are most commonly expressed as integers. 5. A molecule is a tightly bound combination of two or more atoms that acts as a single unit.
Evidence for these postulates Law of Conservation of Mass – Discovered by Antoine Laurent Lavoisier. States that matter can neither be created nor destroyed. There is no detectable change in mass in an ordinary chemical reaction. Lavoisier proved this law by conducting many experiments in which he showed that eth total mass of matter at the end of the experiment was exactly the same as that at the beginning. Law of Constant Composition – Discovered by Joseph Proust. States that any compound is always made up of elements in the same proportion by mass.
Monatomic Elements – Consists of single elements that are not connected to each other. Example – Helium & Neon.
Diatomic Elements (memorize these) – They contain two atoms of the same element per molecule Example – H2, N2, F2, Cl2, Br2, and I2
Polyatomic Elements (memorize these) – Giant clusters of bonded atoms. Example – Ozone= O3; Phosphorus= P4; Diamond has millions of carbon atoms all bonded together in a gigantic cluster.

2.4 – What Are Atoms Made Of?
Protons – Has a positive charge. One proton has a charge of +1. The mass of a proton is 1.6726 x 10^-24g, but this number is so small that it is more convenient to use another unit called amu. 1 amu= 1.6605 x 10^-24g. A proton has a mass of 1.0073 amu. We round to one significant figure to get 1 amu.
Neutrons – Has no charge. Neutrons neither attract nor repel each other or any other particles. The mass of a neutron is slightly greater than that of a proton. 1.6749 x 10^-24g or 1.0087amu. We round this up to have one significant figure, which gives you 1 amu.
Electrons – Has a charge of -1. Is equal in magnitude to the charge on a proton, but opposite in sign. The mass of an electron is approximately 5.4858 x 10^-4amu or 1/1837 that of a proton. It takes about 1837 electrons to equal the mass of one proton.
Subatomic Particles
Charge
Mass (g)
Mass (amu)
Mass (amu); Rounded to one Significant Figure
Location in an Atom
Proton
+1
1.6726 x 10^-24
1.0073
1
In the nucleus
Electron
-1
9.1094 x 10^-28
5.4858 x 10^-4
0.0005
Outside the nucleus
Neutron
0
1.6749 x 10^-24
1.0087
1
In the nucleus

Nucleus – Is the tight cluster of protons and neutrons located in the center of an atom. Electrons are found as a diffuse cloud outside the nucleus.
Atomic Mass Units – When the mass if so small that it is more convenient to use another unit. Also known as amu. Used to tell the mass of protons, electrons, and neutrons.
Mass Number – One way to describe an atom. The sum of the number of protons and neutrons in its nucleus. Mass number (A)= the number of protons + neutrons in the nucleus of an atom.
Atomic Number – The atomic number of an element is the number of protons in its nucleus. Atomic number (Z)= number of protons in the nucleus of an atom.
Isotopes – Atoms with the same number of protons but different numbers of neutrons. Each isotope, contains a different number of neutrons and, therefore, has a different mass number.
Atomic Weight – Is a weighted average of the mass (in amu) of its isotopes found on the earth. To calculate the weighted average of the masses of the isotopes, multiply each atomic mass by its abundance and then add. Example – Magnesium= The natural abundances of the three stable isotopes of magnesium are 78.99% magnesium-24 (23.98504 amu), 10.00% magnesium-25 (24.9858 amu), and 11.01% magnesium-26 (25.9829 amu). (78.99/100 x 23.99 amu) + (10.00/100 x 24.99 amu) + (11.01/100 x 25.98 amu)= 18.95+2.499+2.860= 24.31 amu.

2.5 – What Is the Periodic Table?
Origins – 1860s. Dmitri Mendeleyev produced the first periodic table. This is the same form we use today.
Periods – Horizontal rows. Starting anew row each time he came to an element with properties similar to hydrogen.
Groups – The groups are numbered 1 to 18, without adding letters, beginning on the left. The A group elements (Groups 1A and 2A on the left side of the table and Groups 3A through 8A at the right) are known collectively as main-group elements.
Metals – The majority of elements are metal. Metals are solid at room temperature, except for mercury, shiny, conductors of electricity, ductile, and malleable. In reactions they tend to give up electrons. They also form alloys.
Nonmetals – Are the second class of elements. With the exception of Hydrogen they are located on the right side of the periodic table. There are 19 non metals. They do not conduct electricity; except for graphite.
Metalloids/Semimetals – There are 6 of them. They are boron, silicon, germanium, arsenic, antimony, and tellurium. These elements have some properties of metals and some of nonmetals. Silicon is a semiconductor.
Alkali Metals – All of these are soft enough to be cut with a knife, and their softness increases going down the column. Involves group 1A. They have low melting and boiling points; they decrease going down the column. All alkali metal react with water to form hydrogen gas, and a metal hydroxide with the formula MOH. The violence of their reaction with water increases in going down the column.
Transition Metals – The elements in the B columns. Groups 3 to 12 in the new numbering system. Notice the elements 58 to 71 and 90 to 103 are not included in the main body of the table but rather are shown separately at the bottom theses are known as inner transition elements.
Halogens – Are all color substances. Fluorine (atomic number 9), Chorine (17), Bromine (35), & Iodine (53). All form compounds with sodium that have the general formula NaX; NaCl & NaBr. Only the elements in this column share this property.
Noble Gases – Elements in group 8A. They are gases under normal pressure and temperature, and they form either no compounds or very few compounds. Their melting and boiling point sare very close to one another.

2.6 – How Are Electrons in an Atom Arranged?
Discovery of Neils Bohr – An electron is always moving around the nucleus and so possesses kinetic energy. He discovered only certain values are possible for this energy.
Ground State – Is the lowest possible energy level. If an electron is to have more energy than it has in eth ground state, only certain values are allowed; values in between are not permitted. The electron configuration of the lowest energy state of an atom.
Shells – Electrons in atoms do not move freely in the space around the nucleus, but rather remain confined to the shells. There are 4 different shells.
Shell
Number of Electrons Shall can Hold
Relative Energies of Electrons in Each Shell
4
32
Higher ^
3
18 |
2
8 |
1
2
Lower |

Orbitals – A region of space around a nucleus that can hold a maximum of two electrons.
Electron Configuration – A description of the orbitals of an atom or ion occupied by electrons. Example(s) – 1s, 2s, 2p, 3s, 3p, … Rule1= Orbitals fill in the order of increasing energy from lowest to highest Rule2= Each orbital can hold up to two electrons with spin pairs Rule3= When there is a set of orbitals of equal energy, each orbital becomes half-filled before any of them becomes completely filled.

Orbital Diagrams – We use a box to represent an orbital, an arrow with its head up to represent a single electron, and a pair of arrows with head in opposite directions to represent two electron configurations for elements 1 through 18. Example(s) –

Lewis Dot Structures – Shows the symbol of the element surrounded by a number of dots equal to the number of electrons in the outer (valence) shell of an atom of that element. Example(s) –

2.7 – How are Electron Configurations and Position of the Periodic Table Related?
Compare groups to electron configurations – Elements in the same column have the same ground-state electron configuration in their outer shells. All main group elements (those in A column) have in common the fact that either their s or p orbitals are filled with eight electrons.

2.8 – What is a Periodic Property?
Periodicity – Atomic size & ionization energy. Periodicity is related to the positioning on the Periodic Table.
Atomic Size – The size of an atom is determined by the size of its outermost occupied orbital. Example – The size of a chlorine atom is determined by the size of its three 3p orbitals (3s^23p^5). The simplest way to determine the size of an atom is to determine the distance between bonded nuclei in a sample of the element. A chlorine molecule, for example, has a diameter of 198 pm (pm=picometer; 1pm= 10^-12 meter). The radius of a chlorine atom is thus 99pm, which is one-half of the distance between two bonded chlorine nuclei in Cl2.

Ionization Energy – Is the measure of how difficult it is to remove the most loosely held electron, the higher the ionization energy. Example –