Chemistry Study Guide
Unit 4 – Bonding and Chemical Formulas
Name
Test Deadline
Relevant State Content Standards:
1d –Students know how to use the periodic table to determine the number of electrons
available for bonding.
2. Chemical Bonds –biological, chemical, and physical properties of matter result
from the ability of atoms to form bonds from electrostatic forces between electrons
and protons and between atoms and molecules.
2a – Atoms combine to form molecules by sharing electrons to form covalent or metallic
bonds or by exchanging electrons to form ionic bonds
2b – Chemical bonds between atoms in molecules such as H2, CH4, NH3, H2CCH2, N2,
Cl2 and many large biological molecules are covalent.
2c – Salt crystals, such as NaCl, are repeating patterns of positive and negative ions held
together by electrostatic attraction.
2e – Students know how to draw Lewis dot structures.
2f – Predict the shape of simple molecules and their polarity from Lewis dot structures
2g – Know how electronegativity and ionization energy relate to bond formation
Other Unit Goals:
1 – Know Cations are positive ions; Anions are negative ions
2 – Monatomic Ions contain 1 element; Polyatomic ions contain two or more elements.
3 – Oxidation numbers tell how many electrons are gained or lost when an atom becomes
an ion.
4 – When writing a chemical formula, the total positive charge must equal the total
negative charge.
5 – Know the meaning of the term Diatomic Element, and the 7 diatomic elements.
6 – Be able to name and write the chemical formula for any monatomic/polyatomic ion,
or ionically or covalently bonded compound. See figure 9.20 and 9.22 on pages 277 and
278 for a flow chart on how to name and write formulas.
Text Reference: Chapters 7, 8, and 9 (pages 186 – 285)
Chemistry Study Guide
Unit 4 – Bonding and Chemical Formulas
Name
Test Deadline
Questions: (Please use complete sentences, show work, and answer on a separate sheet
of paper)
1.) What are chemical formulas?
2.) Why is it important to learn formula writing and naming?
3.) Explain the difference between an ionic compound and a molecular (covalent)
compound? Give a specific example of each.
4.) What is electronegativity?
5.) Examine the values of electronegativity on p.177. What is the relationship between
the position of an element in a period and in a family to the amount of electronegativity
that the atom has?
6.) What is a polyatomic ion? A monatomic ion?
7.) What do numbers and letters represent in chemical formulas?
8.) When do you use subscripts and parenthesis in writing chemical formulas?
9.) What is meant by the term Diatomic Element? What are the 7 diatomic elements?
10.) Draw the Lewis electron-dot notations for each atom below:
A) Calcium B) Argon
C) Boron
D) Chlorine
E) Fluorine
Book Questions: pg. 193 (1 – 10); pg. 211 (1-13); pg. 244 (32, 34); pg. 251 (1, 4, 5); pg.
258 (3 – 9); pg. 266 (14 – 19); pg. 270 (20, 23); pg. 285 (1 – 11) – 54 book questions
Chemistry Study Guide
Unit 4 – Bonding and Chemical Formulas
Name
Test Deadline
Supplemental Information:
Compounds may either be ionic or covalent (molecular). Ionic Compounds form by
the transfer of electrons from one atom to another. When at atom gains electrons it
becomes a negative ion or anion. When an atom loses electrons it becomes a positive ion
or cation. It is the attraction between these positive and negative ions that forms the
ionic bond. Atoms at opposite ends of the table form ionic bonds (Typically a metal
bonded to a nonmetal). This is because they have a large difference in electronegativity.
Polyatomic ions are composed of two or more atoms with an overall charge: SO4-2, NH4+1
Covalent compounds are formed by sharing electrons. The result is the formation of a
molecule. The nonmetals tend to form molecules.
Oxidation numbers tell you the charge of an atom when it becomes an ion. The charge
of the Group I elements is +1 because they lose one electron. Group II is +2, Group III is
+3 and Group IV is +4 or -4. Group V is -3 because it gains three electrons. Group VI is
-2 and Group VII is -1. The Noble Gases in Group VIII have a zero oxidation number
because they cannot accept or lose electrons. The transition metals and inner transition
metals may have more than one oxidation number.
Lewis Electron-Dot Notation:
Electron-dot notation is another way to show the electrons in an atom; however, it only
represents the outermost energy level. Dots are used to represent electrons. They are
placed in a specific order around the symbol for the element.
6 3
4
7
X
2
1
5
8
For example, take the element nitrogen. It electron configuration is 1s22s22p3.
The second energy level is the highest level containing electrons. It contains a total of
five electrons. The electron-dot notation for nitrogen is:
N