Chemistry
Ms. Boon
Unit 2: Chemical Bonds Study Guide
Name: ____________________
Period: ____ Date: ______
Standard 1d: I know how to use the periodic table to determine the number of electrons available for
bonding.
Main Ideas:
 Protons determine the physical properties of an atom. The number of protons tells us what element an
atom is.
 Electrons determine the chemical properties of an atom. Electrons are involved in chemical reactions
and chemical bonding.
 A period on the periodic table is a horizontal row. The periods represent energy levels.
 Electrons are arranged in energy levels and orbitals around the nucleus. The electron configuration of
an atom tells us approximately where the electrons are located.
 Electrons fill the lowest energy levels first.
 The electrons in the outer or highest energy level that are available for bonding are called valence
electrons.
 A group on the periodic table is a vertical column. Elements in the same group generally have the same
number of valence electrons (exception: the transition metals).
 Atoms tend to lose or gain electrons, forming ions, in order to have a full outer energy level (shell) of
electrons. This is called the octet rule because a full outer shell is usually 8 electrons (exception 1st
energy level with only the 1s2 orbital – holds 2 electrons.)
 The noble gases are very stable because they already have a full outer shell of electrons.
 When an atom forms an ion according to the octet rule, that ion will have the same number of electrons
as the closest noble gas.
Skills:
 Use periodic table to determine the period, group, electron configuration, and number of valence
electrons of main group elements (group 1, 2, 13, 14, 15, 16, 17, and 18).
 Use octet rule to determine what ion an atom will form. Explain the octet rule and relate to stability of
noble gases and reactivity of other main group elements like the alkali metals, halogens, and alkaline
earth metals.
Note: The electron configuration pattern will be written on the board. (1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6)
Practice: Complete the valence electron and electron configuration review WS.
Standard 2c: I know salt crystals, like NaCl (table salt), are repeating patterns of positive and
negative ions held together by electrostatic (attraction of oppositely charged particles) attractions.
Main Ideas:
 An electron is a negatively charged subatomic particle
 An ion is an atom that has lost or gained one or more electrons.
 Metals form positive cations by losing electrons.
 Nonmetals form negative anions by gaining electrons.
 Octet rule: generally atoms gain or lose electrons in order to have a full outer shell of electrons
like a noble gas.
 Ionic bonds form between metals and nonmetals
 Ionic bonds are held together by the electrostatic attraction between negative anions and
positive cations. (Opposites attract!)


To form an ionic bond, a metal atom transfers an electron to a nonmetal atom. The transfer
occurs in a ratio that results in a neutral compound.
Ionic compounds form a crystal lattice structure which is a highly organized arrangement of
anions and cations. All anions will be surrounded by cations and vice versa in order to form a
neutral compound.
Skills:
 Write ionic compound formulas (use Criss cross rule)
 Write ionic compound names (remember what roman numerals mean and that anions have
“ide” at the end).
Practice: Ions review worksheet. Behind the electron configuration review distributed
Monday/Tuesday.
Standard 2a: Students know atoms combine to form molecules by sharing electrons to form covalent
or metallic bonds or by exchanging electrons to form ionic bonds.
Main Ideas:
 A chemical bond is an attraction between atoms caused by rearranging their valence electrons
in some way.
 The types of chemical bonds we have studied are ionic bonds and covalent bonds. A third type
of chemical bond is the metallic bond. Electrons are shared in metallic bonds, as in covalent
bonds.
 Bond strength or bond energy is the energy required to break a covalent bond. The bond length
is the distance between the nuclei of bonded atoms. As bond length increases, bond strength
decreases. In other words, the shorter the bond length, the stronger the bond.
Ionic bonds
Covalent bonds
Ionic bonds form ionic compounds.
Covalent bonds form covalent or molecular
compounds.
In an ionic bond, electrons are transferred from
a metal to a nonmetal atom.
In covalent bonding, valence electrons are share by
two atoms.
Ionic bonds form between metals and
nonmetals
The structure of an ionic compound is a crystal
lattice. In the crystal lattice, all the positive
cations are surrounded by negative anions and
vice versa. All this attraction makes a strong
structure.
Covalent bonds form between nonmetals
Ionic compounds are solids (unless heated to
very high temperatures)
Covalent compounds can be solid, liquid, or gas.
The structure of a covalent compound is a molecule.
Only the atoms that share electrons are strongly
attracted to one another. The individual covalent
bonds are strong, but the overall strength of the
bonds in the compound are weak compared to ionic
compounds.
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Ionic compounds have very high melting and
boiling points.
Covalent compounds have lower melting and boiling
points.
Ionic compounds conduct electricity when they
are dissolved in a solution.
Covalent compounds do not conduct electricity.
Skills: Use knowledge of what elements are metals and what elements are nonmetals to distinguish
between ionic and covalent compounds.
Practice: Review problems pp. 216 # 11, 12, 19, pp. 217 #39, pp. 218 #42, (challenge #46)
Standard 2b: Students know chemical bonds between atoms in molecules are covalent bonds.
Main Ideas:
 A molecule consists of two or more atoms covalently bonded together.
 Molecules have their own physical properties, different from the individual atoms that make the
molecule. (Example: water is made of hydrogen and oxygen and the properties of water are
very different from the properties of hydrogen and oxygen separately.)
 Covalent bonds usually form between nonmetals.
 Macromolecules and other biological compounds are made of molecules formed by covalent
bonds.
 Because carbon has 4 valence electrons, every carbon atom can form 4 covalent bonds. This
makes carbon great for forming very large molecules. Carbon atoms are usually the “central
atoms” or form the skeleton of large biological molecules.
 A single covalent bond is the sharing of a pair of electrons between two atoms.
 A double covalent bond is the sharing of two pairs of electrons between two atoms.
 A triple covalent bond is the sharing of three pairs of electrons between two atoms.
 An unshared pair of electrons is a pair of valence electrons that is not involved in a covalent
bond.
 In a nonpolar covalent bond, electrons are shared equally.
 In a polar covalent bond, electrons are not shared equally.
 Polarity depends on the difference in electronegativity between the atoms sharing electrons.
(see pp. 194 for electronegativity data).
 Atoms form covalent bonds in order to fill their outer energy level with electrons (8 electrons).
This is the octet rule (or the duet rule for hydrogen – 2 electrons).
 Forming a covalent bond that meets the octet rule makes the molecule more stable than the
individual atoms.
Skills:


Understand when a single, double, or triple bond will form.
Explain why nonmetals and nonmetals form covalent bonds rather than ionic bonds.
Practice: pp. 198 #4, pp. 217 #40, pp. 218 #48
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Standard 2e: I can draw Lewis Dot structures of molecules
Main ideas:

Lewis dot structures represent valence electrons and covalent bonds.

Atoms in the same group on the periodic table have the same number of valence electrons.

A single bond, sharing a pair of electrons between two atoms, is represented by a single line.

A double bond, sharing two pairs of electrons between two atoms, is represented by two lines.

A triple bond, sharing three pairs of electrons between two atoms, is represented by three lines.

Atoms form covalent bonds in order to satisfy the octet rule.
Skills: Draw Lewis dot structures of molecules. Remember to include unshared pairs.
Practice: Review the classwork and HW worksheets on drawing molecules.
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