Chapter 1

• All matter is composed of the same building blocks called atoms.
• two main components of an atom:

  1. The nucleus contains positively charged protons and uncharged neutrons. Most of the mass of the atom is contained in the nucleus.
  2. The electron cloud is composed of negatively charged electrons. The electron cloud comprises most of the volume of the atom.
• In a neutral atom, the number of protons in the nucleus equals the number of electrons. This quantity, called the atomic number, is unique to a particular element.
• In addition to neutral atoms, we will also encounter charged ions.

  • A cation is positively charged and has fewer electrons than its neutral form.
  • An anion is negatively charged and has more electrons than its neutral form
• Isotopes are two atoms of the same element having a different number of neutrons. Isotopes have different mass numbers.
• The mass number of an atom is the total number of protons and neutrons in the nucleus.
• The atomic weight is the weighted average of the mass of all isotopes of a particular element, reported in atomic mass units (amu).
• A row in the periodic table is also called a period, and a column is also called a group.
• Long ago it was realized that groups of elements have similar properties, and that these atoms could be arranged in a schematic way called the periodic table.
• The periodic table is composed of rows and columns.

  • Elements in the same row are similar in size.
  • Elements in the same column have similar electronic and chemical properties
• Each shell contains a certain number of subshells called orbitals. An orbital is a region of space that is high in electron density. There are four different kinds of orbitals, called s, p, d, and f. The first shell has only one orbital, called an s orbital. The second shell has two kinds of orbitals, s and p, and so on. Each type of orbital occupies a certain space and has a particular shape.
• An s orbital has a sphere of electron density. It is lower in energy than other orbitals of the same shell, because electrons are kept close to the positively charged nucleus. An s orbital is fi lled with electrons before a p orbital in the same shell.
• A p orbital has a dumbbell shape. It contains a node of electron density at the nucleus. A node means there is no electron density in this region. A p orbital is higher in energy than an s orbital (in the same shell) because its electron density is farther away from the nucleus. A p orbital is filled with electrons only after an s orbital of the same shell is full.
• Each orbital can have a maximum of two electrons.
• Every element in the second row has a filled first shell of electrons. Thus, all second-row elements have a 1s^2 configuration. These electrons in the inner shell of orbitals are called core electrons, and are not usually involved in the chemistry of a particular element.
• Each element in the second row of the periodic table has four orbitals available to accept additional electrons:

  • one 2s orbital, the s orbital in the second shell
  • three 2p orbitals, all dumbbell-shaped and perpendicular to each other along the x, y, and z axes
• Because each of the four orbitals in the second shell can hold two electrons, there is a maximum capacity of eight electrons for elements in the second row
• The outermost electrons are called valence electrons.
• The valence electrons are more loosely held than the electrons closer to the nucleus, and as such, they participate in chemical reactions. The group number of a second-row element reveals its number of valence electrons.
• Bonding is the joining of two atoms in a stable arrangement.
• Bonding is a favorable process because it always leads to lowered energy and increased stability.
• Joining two or more elements forms compounds.
• One general rule governs the bonding process.

  • Through bonding, atoms attain a complete outer shell of valence electrons.
• Alternatively, because the noble gases in column 8A of the periodic table are especially stable as atoms having a filled shell of valence electrons, the general rule can be restated.

  • Through bonding, atoms attain a stable noble gas configuration of electrons.
• A first-row element like hydrogen can accommodate two electrons around it.
• A second-row element is most stable with eight valence electrons around it like neon. Elements that behave in this manner are said to follow the octet rule.
• two different kinds of bonding: ionic bonding and covalent bonding.
• Ionic bonds result from the transfer of electrons from one element to another.
• Covalent bonds result from the sharing of electrons between two nuclei.
• An ionic bond generally occurs when elements on the far left side of the periodic table combine with elements on the far right side, ignoring the noble gases, which form bonds only rarely. The resulting ions are held together by extremely strong electrostatic interactions.
• The transfer of electrons forms stable salts composed of cations and anions.
• A compound may have either ionic or covalent bonds. A molecule has only covalent bonds.
• covalent bonding, occurs with elements like carbon in the middle of the periodic table, which would otherwise have to gain or lose several electrons to form an ion with a complete valence shell.
• A covalent bond is a two-electron bond, and a compound with covalent bonds is called a molecule.
• Second-row elements can have no more than eight valence electrons around them.
• For neutral molecules, two consequences result.

  • Atoms with one, two, three, or four valence electrons form one, two, three, or four bonds, respectively, in neutral molecules.
  • Atoms with five or more valence electrons form enough bonds to give an octet. This results in the following simple equation: predicted number of bonds = 8 – number of valence electrons
• Nonbonded pair of electrons = unshared pair of electrons = lone pair
• when second-row elements form fewer than four bonds their octets consist of both bonding (shared) electrons and nonbonding (unshared) electrons. Unshared electrons are also called lone pairs.
• Lewis structures are electron dot representations for molecules.
• There are three general rules for drawing Lewis structures.

  1. Draw only the valence electrons.
  2. Give every second-row element no more than eight electrons.
  3. Give each hydrogen two electrons.