Bonding/Moles Unit
Table
of
Contents
Chapter
7
-‐
Ionic
Bonding
. ...........................................................................................................................................
3
7.1
Ions
. ..........................................................................................................................................................................................
4
7.2
Ionic
Bonds
and
Ionic
Compounds .
................................................................................................................................
6
7.1/7.2
Homework
.........................................................................................................................................................................................
8
Chapter
7
Practice
Quiz .
...............................................................................................................................................
9
Chapter
7
Study
guides .
.............................................................................................................................................. 1
1
7.1:
Ions
–
Study
Guide
. ...........................................................................................................................................................
1 1
7.2:
IONIC
BONDS
AND
IONIC
COMPOUNDS
–
Study
guide
. .........................................................................................
1 2
Chapter
8
-‐
Covalent
Bonding
.................................................................................................................................. 1
3
8.1
-‐
Molecular
Compounds
. ..................................................................................................................................................
1 4
8.2
-‐
The
Nature
of
Covalent
Bonding
................................................................................................................................
1 5
Drawing
Lewis
Structures
for
Covalent
Compounds
Worksheet
.............................................................................
1 8
8.1/8.2
Homework
(I
have
to
check
#s
20-‐23
before
scanning) .
............................................................................................
2 0
8.4
Polar
Bonds
and
Molecules
............................................................................................................................................
2 2
8.4
Homework
(I
must
check
#s
11
and
12
before
you
scan)
..................................................................................................
2 4
Chapter 8 Practice Quiz .
........................................................................................................................................... 2
5
Chapter
8
Study
Guides
. ............................................................................................................................................. 2
7
8.1:
MOLECULAR
COMPOUNDS
–
Study
guide
.................................................................................................................
2 7
8.2:
THE
NATURE
OF
COVALENT
BONDING
–
Study
guide
..........................................................................................
2 9
8.4:
POLAR
BONDS
AND
MOLECULES
–
Study
guide
. .....................................................................................................
3 0
Chapter
10
-‐
Chemical
Quantities
. .......................................................................................................................... 3
2
10.1
Chemical
Measurements .
..............................................................................................................................................
3 3
10-‐1
Worksheet
–
What
is
a
mole?
. .....................................................................................................................................
3 4
10-‐2
Mole
Conversions .
...........................................................................................................................................................
3 6
10.1/10.2
Homework
(you
must
show
work
for
credit)
. ...........................................................................................................
3 8
10-‐3
Empirical
and
Molecular
Formulas
..........................................................................................................................
4 0
10.3
Homework
............................................................................................................................................................................................
4 1
Chapter
10
Practice
Quiz
. .......................................................................................................................................... 4
2
Chapter
10
Study
Guides .
........................................................................................................................................... 4
4
10.1:
The
Mole:
A
Measurement
of
Matter
–
Study
guide
. ...........................................................................................
4 4
10.2:
Mole-‐Mass
and
Mole-‐Volume
Relationships
–
Study
guide .
.............................................................................
4 6
10.3:
Percent
Composition
and
Chemical
Formulas
–
Study
guide
. .........................................................................
4 8
Ionic
or
Covalent
Bonding
Lab .
................................................................................................................................ 5
0
Dot
Structures
and
Molecular
Model
Building
Lab .
.......................................................................................... 5
2
Moles,
atoms,
grams
oh
my!
Lab
............................................................................................................................. 5
5
Mole
conversions
summary
sheet
. ......................................................................................................................... 5
7
Unit
Test
Setup
. ............................................................................................................................................................. 5
8
Unit
Review
Materials
................................................................................................................................................ 5
9
Chapter
7
vocabulary
review
. ...............................................................................................................................................
5 9
Chapter
10
vocabulary
review
.............................................................................................................................................
6 0
Unit
Review
–
Chapters
7,
8,
and
10 .
...................................................................................................................................
6 1
2
Chapter
7
-‐
Ionic
Bonding
7.1 – Ions
Objectives
• Determine the number of valence electrons in an atom of a representative element
• Explain the octet rule
• Describe how cations form
• Explain how anions form
Vocabulary • octet rule
• halide ions
• valence electrons
• electron dot structures
7.2 – Ionic Bonds and Ionic Compounds
Objectives
• Explain the electrical charge of an ionic compound
• Describe three properties of ionic compounds
Vocabulary • chemical formula
• formula unit
• ionic compounds
• ionic bonds
3
7.1
Ions
Lesson Summary - After reading Lesson 7.1, fill in the following statements.
Valence Electrons
Valence electrons are the electrons in the ________________ occupied energy level and are involved in ion
formation.
For a representative element, the group number equals the number of ______________ electrons
the atom contains.
An electron dot structure shows the ________________ of the element and its valence electrons.
Atoms tend to _________ or _________ the number of electrons that will provide the atom with a noble gas
electron configuration.
Formation of Cations
Cations are ________________ charged ions formed when an atom ___________ one or more valence electrons.
Atoms and the cations formed from them have __________________ properties.
Elements in Group 1 form cations with a charge of ___, and those in Group 2 form cations with a charge of ___.
Many ________________ metals form more than one cation and do not follow the ___________ rule.
Formation of Anions
Anions are ___________________ charged ions formed when an atom ___________ one or more valence electrons.
Commonly, the name of an anion ________ in -ide.
Anions form from ___________________ elements.
The anions formed from halogens are known as _______________.
Valence Electrons
1. What are valence electrons?
2. The valence electrons largely determine the of an element and are usually the
only electrons used in .
3. Is the following sentence true or false? The group number of a representative element in the periodic table is
related to the number of valence electrons it has. _____________
4. What is an electron dot structure?
5. Draw the electron dot structure for each of the following atoms.
Ar Ca I
a. argon b. calcium c. iodine
6. What is the octet rule?
7. Metallic atoms tend to valence electrons to produce a positively charged ion. Most
nonmetallic atoms achieve a complete octet by gaining or ____________________ electrons.
4
Formation of Cations
8. Using your periodic table, write the electron configurations for these metals, and circle the electrons lost when
each metal forms a cation.
a. Mg b. Al _____________________ c. K _________________________
Match the noble gas with its electron configuration.
9. argon a. 1s2
b. [He]2s22p6
10. helium c. [Ne]3s23p6
11. neon d. [Ar]3d104s24p6
12. krypton
13. What is the electron configuration called that has 8 electrons in the outer energy level and all of the orbitals
filled?
14. Write the electron configuration for zinc.
15. Fill in the electron configuration diagram for the copper(I) ion.
Energy level
Copper atom Copper (I) ion
Cu Cu+
Formation of Anions
16. Atoms of most nonmetallic elements achieve noble-gas electron configurations by
gaining electrons to become , or negatively charged ions.
17. What property of nonmetallic elements makes them more likely to gain electrons than lose electrons?
18. True or false? Elements of the halogen family lose one electron to become halide ions. ____________
19. How many electrons will each element gain in forming an ion?
a. nitrogen b. oxygen c. sulfur d. bromine
20. Write the symbol and electron configuration for each ion from Question 19, and name the noble gas with the
same configuration.
a. nitride
b. oxide
c. sulfide
d. bromide
5
7.2
Ionic
Bonds
and
Ionic
Compounds
Lesson Summary - After reading the section, fill in the following statements.
Formation of Ionic Compounds An ionic compound is made up of __________ and cations and has an
overall charge of __________.
The __________________ attraction between an anion and a cation is an ionic bond.
The representative unit of an ionic compound is its __________________________.
A formula unit of an ionic compound shows the ions in their lowest, ______________________ ratio.
Properties of Ionic Compounds Ionic compounds have characteristic properties that distinguish them from
other substances.
Most ionic compounds are ___________________ solids at room temperature.
In general, ionic compounds have __________ melting points because the ions have a strong attraction for one
another.
Ionic compounds conduct an electric current when ________________ or in an ______________ solution
because the ions are then free to _____________.
Formation of Ionic Compounds
1. What is an ionic bond?
2. In an ionic compound, the charges of the and must balance to
produce an electrically
substance.
3. Complete the electron dot structures below to show how beryllium fluoride (BeF2) is formed. Use the diagram
on page 203 as a model.
4. Why do beryllium and fluorine combine in a 1:2 ratio?
5. A chemical formula shows the types and of atoms in the smallest representative unit of a
substance.
6. List the numbers and types of atoms represented by these chemical formulas.
a. Fe2O3
b. KMnO4
c. CH3
d. NH4NO3
6
7. What is a formula unit?
8. Explain why the ratio of magnesium ions to chloride ions in MgCl2 is 1:2.
9. Describe the structure of ionic compounds.
Properties of Ionic Compounds
10. Most ionic compounds are at room temperature.
11. T or false? Ionic compounds generally have low melting points. __________
12. Circle the letter of each statement that is true about ionic compounds.
a. When dissolved in water, ionic compounds can conduct electricity.
b. When melted, ionic compounds do not conduct electricity.
c. Ionic compounds have very unstable structures.
d. Ionic compounds are electrically neutral.
7
7.1/7.2
Homework
__1. The Lewis dot symbol for a lead atom is:
A. B. C. D. E.
__2. The Lewis dot symbol for the S2- ion is:
2- E.
A. B. C. S2- D.
__3. The Lewis dot symbol for the chloride ion is:
A. B. C.
D. E. __12. Which of the following elements forms an ion with a 1– charge?
__4. The Lewis dot symbol for the calcium ion is: A. F B. O C. K D. Na
A. B. C. D. Ca2+ __13. What is the formula of the ion formed when tin achieves a noble
E. Ca
gas electron configuration?
A. Sn+1 B. Sn+2 C. Sn+3 D. Sn+4
__5. Select the element whose Lewis symbol is correct. __14. What is the formula of the ion formed when phosphorus
achieves a noble gas electron configuration?
A. P+1 B. P+3 C. P-3 D. P-1
__6. Select the element whose Lewis symbol is __15. How does oxygen obey the octet rule when reacting to form
correct. compounds?
A. It gains electrons.
B. It gives up electrons.
C. It does not change its number of electrons.
D. Oxygen does not obey the octet rule.
__7. Select the correct formula for a compound __16. Which of the following occurs in an ionic bond?
A. Oppositely charged ions attract.
formed from calcium and chlorine. B. Two atoms share two electrons.
C. Two atoms share more than two electrons.
A. CaCl B. CaCl2 C. Ca2Cl D. Like-charged ions attract.
D. Ca2Cl2 E. CaCl3
__8. How does calcium obey the octet rule when __17. Which of the following is true about an ionic compound?
reacting to form compounds? A. The chemical formula shows the atoms in a molecule.
A. It gains electrons. B. The formula unit gives the number of each type of ions in a crystal.
B. It gives up electrons. C. It is composed of anions and cations and yet it is electrically
C. It does not change its number of electrons. neutral.
D. Calcium does not obey the octet rule. D. The chemical formula shows the ions in a molecule.
__9. What is the electron configuration of the calcium __18. How many valence electrons are transferred from the nitrogen
ion?
A. 1s22s22p63s23p6 atom to potassium in the formation of the compound potassium nitride?
B. 1s22s22623s23p64s2
C. 1s22s22p63s23p64s24p6 A. 0 B. 1 C. 2 D. 3
D. 1s22s22p63s23p63d2
__19. How many valence electrons are transferred from the calcium
atom to each iodine atom in the formation of the compound calcium
__10. What is the charge on the strontium ion? iodide?
A. 2– B. 1– A. 0 B. 1 C. 2 D. 3
C. 1+ D. 2+ __20. What is the formula unit of aluminum oxide?
__11. What is the formula of the ion formed when A. AlO B. Al2O C. Al3O2 D. Al2O3
potassium achieves noble-gas electron configuration? __21. Ionic compounds are normally in which physical state at room
A. K+1 B. K+2
C. K-1 D. K-2 temperature?
A. solid B. liquid C. gas D. plasma
8
Chapter
7
Practice
Quiz
Matching: Match the NUMBER of valence electrons on the right with the element on the left.
NOTE: Some may be used more than once and some may not be used at all.
__1. Phosphorus 1
__2. Vanadium 2
__3. Oxygen 3
__4. Chromium 4
__5. Chlorine 5
__6. Helium 6
__7. Gallium 7
__8. Carbon 8
__9. What is the name given to the electrons in the highest occupied energy level of an atom?
A. orbital electrons B. valence electrons C. anions D. cations
__10. How does calcium obey the octet rule when reacting to form compounds?
A. It gains electrons. B. It gives up electrons.
C. It does not change its number of electrons. D. Calcium does not obey the octet rule.
__11. What is the electron configuration of the calcium ion?
A. 1s22s22p63s23p6 B. 1s22s22p63s23p34s2
C. 1s22s22p63s23p54s1 D. 1s22s22p63s2
__12. What is the electron configuration of the gallium ion?
A. 1s22s22p63s23p6 B. 1s22s22p63s23p54s1
C. s22s22p63s23p54s24p6 D. 1s22s22p63s23p63d10
__13. What is the charge on the strontium ion? 2– 1– 1+ 2+
__14. The octet rule states that, in chemical compounds, atoms tend to have ____.
A. the electron configuration of a noble gas B. more protons than electrons
C. eight electrons in their principal energy level D. more electrons than protons
__15. How many electrons does silver have to give up in order to achieve a pseudo-noble-gas electron
configuration? 1234
__16. How many electrons does barium have to give up to achieve a noble-gas electron configuration?
1234
__17. What is the formula of the ion formed when potassium achieves noble-gas electron configuration?
A. K2+ B. K1+ C. K1- D. K2-
__18. Which of the following elements forms an ion with a 1– charge?
A. fluorine B. hydrogen C. potassium D. sodium
__19. What is the formula of the ion formed when tin achieves a stable electron configuration?
A. Sn4+ B. Sn3+ C. Sn2- D. Sn4-
9
__20. What is the formula of the ion formed when cadmium achieves a pseudo-noble-gas electron
configuration?
A. Cd3+ B. Cd2+ C. Cd1+ D. Cd2-
__21. How many electrons does nitrogen gain in order to achieve a noble-gas electron configuration?
1234
__22. What is the formula of the ion formed when phosphorus achieves a noble-gas electron configuration?
A. P3+ B. P2+ C. P2- D. P3-
__23. How does oxygen obey the octet rule when reacting to form compounds?
A. It gains electrons. B. It gives up electrons.
C. It does not change its number of electrons. D. Oxygen does not obey the octet rule.
__24. The electron configuration of a fluoride ion, F , is _______.
A. 1s22s22p5
B. the same as that of a neon atom
C. 1s22s22p63s1 D. the same as that of a potassium ion
__25. What is the electron configuration of the iodide ion?
A. 1s22s22p63s23p63d104s24p64d105s25p6 B. 1s22s22p63s23p63d104s24p64d10
C. 1s22s22p63s23p63d104s24p64d105s2 D. 1s22s22p63s23p63d104s24p6
__26. Which of the following occurs in an ionic bond?
A. Oppositely charged ions attract. B. Two atoms share two electrons.
C. Two atoms share more than two electrons. D. Like-charged ions attract.
__27. A compound composed of cations and anions is called a(n) _______.
A. diatomic molecule B. polar compound
C. covalent molecule D. ionic compound
10
Chapter
7
Study
guides
7.1:
Ions
–
Study
Guide
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms that are introduced in this section. Each blank
can be completed with a term, short phrase, or number.
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
11. The chlorine atom gains seven electrons when it becomes an ion.
12. The chemical properties of an element are largely determined by the number of valence electrons the element has.
13. Atoms acquire the stable electron structure of a noble gas by losing electrons.
14. An atom of an element in Group 1A has seven valence electrons.
15. Among the Group 1A and 2A elements, the group number of each element is equal to the number of valence electrons in
an atom of that element.
16. Sulfur and magnesium both have two valence electrons.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A Column B
17. electron dot structure a. ions that are produced when halogens gain electrons
b. a depiction of valence electrons around the symbol of an
18. valence electron
19. octet rule element
20. cations c. has the electron configuration of argon
d. an electron in the highest occupied energy level of an
21. anions
22. halide ions element’s atom
23. chloride ion e. Atoms in compounds tend to have the electron
configuration of a noble gas.
f. atoms or groups of atoms with a negative charge
g. atoms or groups of atoms with a positive charg
11
7.2:
IONIC
BONDS
AND
IONIC
COMPOUNDS
–
Study
guide
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms that are introduced in this section. Each blank
can be completed with a term, short phrase, or number.
Anions and cations attract one another by means of ____1___. 1.
The forces of attraction that hold ____2___ charged ions together in 2.
ionic compounds are called ____3___. Although they are composed 3.
of ions, ionic compounds are electrically ____4___. The lowest 4.
whole-number ratio of ions in an ionic compound is called a 5.
___5____. 6.
7.
Nearly all ionic compounds are solid ___6____ at room 8.
temperature. Ionic compounds in general have very ___7____ 9.
melting temperatures. This is because the ____8___ attractive forces 10.
between the ions result in a very ___9____ structure. Ionic
compounds conduct an electric current when in the ___1_0___ state or
dissolved in water.
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
11. During the formation of the compound NaCl, one electron is transferred from a sodium atom to a
chlorine atom.
12. The coordination number of an ion is the number of ions of positive charge that surround the ion in a
crystal.
13. The coordination number of the ion Na+ in NaCl is 6.
14. In forming an ionic compound, an atom of an element gains electrons.
15. Ionic compounds cannot conduct electricity if they are dissolved in water.
Part C Matching
Match each description in Column B to the correct term in Column A
Column A Column B
16. ionic compounds a. compounds composed of cations and anions
17. ionic bonds
18. chemical formula b. the electrostatic forces of attraction binding oppositely charged
ions together
19. formula unit
c. shows the kinds and numbers of atoms in the smallest
representative unit of a substance
d. lowest whole-number ratio of ions in an ionic compound
12
Chapter
8
-‐
Covalent
Bonding
8.1 – Molecular Compounds
Objectives
• Distinguish molecular compounds from ionic compounds
• Identify the information a molecular formula provides
Vocabulary • diatomic molecule • molecular formula
• molecular compound
• covalent bond
• molecule
8.2 – The Nature of Covalent Bonding
Objectives
• State a rule that usually tells how many electrons are shared to form a covalent bond
• Describe how electron dot formulas are used
• Predict when two atoms are likely to be joined by a double or a triple covalent bond
• Distinguish between a single covalent bond and other covalent bonds
• Identify some exceptions to the octet rule
Vocabulary • double covalent bonds
• triple covalent bonds
• single covalent bond
• structural formulas
• unshared pairs
8.4 – Polar Bonds and Molecules
Objectives
• Describe how electronegativity values determine the charge distribution in a polar bond
• Describe what happens to polar molecules when placed between oppositely charged metal
plates
Vocabulary • dipole
• polar bond
• nonpolar covalent bond
• polar covalent bond
• polar molecule
13
8.1
-‐
Molecular
Compounds
Molecules and Molecular Compounds
Sharing Electrons
• Ionic bonds form when the combining atoms ___________ or ____________
electrons.
• Another way that atoms can combine is by ______________ electrons.
• Atoms held together by sharing electrons are joined by a ___________________,
forming a neutral group of atoms called a ___________________.
Oxygen gas consists of oxygen molecules; each molecule consists of _____
covalently bonded oxygen atoms.
O2 is an example of a ____________________- a molecule that contains two atoms.
Other diatomic elements are: H2, N2, Cl2, Br2, I2, and F2 (mark these on your periodic
table)
• A compound composed of molecules is called a _________________
compound. Water (H2O) is an example of a molecular compound.
Representing Molecules
A molecular formula is the _________________ formula of a molecular compound.
A molecular formula shows how many __________ of each element a substance
contains.
The molecular formula of ___________ is H2O.
Butane is commonly used in ______________ and household
torches.
The molecular formula for butane is __________.
Molecular formula shows _______ and numbers of atoms only (NH3)
Structural formula shows ____________ and nonbonding electrons
- give insight to shape of molecule
Perspective drawing - shows bonds and ___ shape (closed arrows N H
indicate _______ coming toward you and open arrows show bonds H
going _______ from you)
Ball-and-stick - show 3D shape, but bonds are ___________________ H
Space-filling - show 3D shape more _________________, but can’t N
easily see the bonds (due to _________________ of electron clouds)
HH H
14
Comparing Ionic and Covalent Compounds
Property Ionic compounds Covalent Compounds
electrons in bond
particles
elements
conductivity
(molten or in solution)
state at room temp.
type of bond
solubility in water
melting point
examples
8.2
-‐
The
Nature
of
Covalent
Bonding
The Octet Rule in Covalent Bonding
Single Covalent Bonds
• The hydrogen atoms in a hydrogen molecule are held together mainly by the
_________________ of the shared electrons to the positive __________.
• Two atoms held together by sharing one pair of electrons are joined by a
_______________________________.
• An electron dot structure such as H:H (or H—H) represents the ___________ pair
of electrons of the covalent bond by two dots (or a dash —).
• The halogens also form single covalent bonds in their _______________
molecules. Fluorine is one example.
• By sharing electrons and forming a single covalent bond, two
fluorine atoms each achieve the electron configuration of
__________.
15
• A pair of valence electrons that is not shared between atoms is called an unshared
pair, also known as a ________ pair or a ___________________ pair.
As you can see in the electron dot structures below, the oxygen atom in water has
____ unshared pairs of valence electrons.
Methane contains ________ single covalent bonds.
The carbon atom has _____ valence electrons and needs four _______ valence
electrons to attain a ________________ configuration.
C
HH
H
Double and Triple Covalent Bonds
Atoms form ____________ or triple covalent bonds if they can attain a noble gas
structure by sharing two or ____________ pairs of electrons.
A ______________________________ is a bond that involves two shared pairs of
electrons. CO2 has two double bonds.
Similarly, a bond formed by sharing _________ pairs of electrons is a triple
covalent bond. N2 has a __________ bond.
Exceptions to the Octet Rule
The octet rule cannot be satisfied in molecules whose total number of valence
electrons is an _____ number. There are also molecules in which an atom has
________, or more, than a complete _________ of valence electrons.
Two plausible electron dot structures can be drawn for the _______ molecule, which
has a total of _____________ valence electrons.
16
N
OO
It is ___________________ to draw an electron dot structure for NO2 that satisfies the
octet rule for all atoms, yet NO2 does exist as a _____________ molecule.
Some molecules with an ________ number of valence electrons, such as some
compounds of __________, also fail to follow the octet rule. (BCl3)
A few atoms, especially ____________ and sulfur, expand the octet to ______
or twelve electrons.
Sulfur hexafluoride (______) and phosphorus pentachloride (_____) are examples.
17
Drawing
Lewis
Structures
for
Covalent
Compounds
Worksheet
1) Use the octet rule to figure out the maximum electrons each atom in the molecule should have,
and add them
up. Every element except H gets 8, and H gets 2 (of course).
2) Count the total valence electrons for the molecule: To do this, draw each atom with the proper
number of dots (valence electrons) and add them up.
3) Subtract the valence electrons from maximum electrons: Or, in other words, subtract the number
you found in #2 above from the number you found in #1 above. The answer you get will be equal to
the number of bonding electrons in the molecule.
4) Divide the number of bonding electrons (#3) by two: Remember, because every bond has two
electrons, the number of bonds in the molecule will be equal to the number of bonding electrons
divided by two.
5) Draw an arrangement of the atoms for the molecule that contains the number of bonds you found
in #4 above: Some handy rules to remember are these:
- Hydrogen and the halogens bond once.
- The oxygen family elements bond twice.
- The nitrogen family elements bond three times. So does boron.
- The carbon family elements bond four times.
A good thing to do is to bond all the atoms together by single bonds, and then add the multiple bonds
until the rules above are followed (and the number of bonds you calculated are complete).
6) Find the number of lone pair (nonbonding) electrons by subtracting the bonding electrons (#3
above) from the valence electrons (#2 above). Arrange these around the atoms until all of them
satisfy the octet rule: Remember, ALL elements EXCEPT hydrogen want eight electrons around
them, total. Hydrogen only wants two electrons.
Example: NBr3
1) The max. # of electrons is equal to 32 (8×4 = 32)
1) Max e- 32
2) The # valence electrons is 5 for N, and 3×7 = 21 for the 3 Br atoms; 5 + 21 = 2) Valence e- 26
26 valence electrons 3) Bonding e- 6
3
3) The # of bonding electrons is equal to the maximum electrons minus the 20
10
valence electrons, (32 – 26 = 6). 4) Bonds
4) The # of bonds is equal to the # of bonding electrons divided by two (6/2 = 5) Nonbonding e-
3), because there are two electrons per bond. As a result, the # of bonds is 3. 6) Nonbonding pairs
5) If we arrange the molecule so that the atoms are held together by three bonds
(3 single bonds).
6) The number of nonbonding e- = the # of valence electrons minus the # of bonding electrons, 26 -
6 = 20.
●● ●● ●●
Br N Br● ●
●● ●
● ●●
Br● ●
● ●18● ●
Write Lewis structures for the following covalent compounds.
SiO2
C2H2
H2O2
HCN
C2H4
CH2F2
C2H5OH
(NH2)2CO
19
8.1/8.2
Homework
(I
have
to
check
#s
20-‐23
before
scanning)
1-10 - Matching: Match the term with the description
__1. molecular formula A. indicate bonds coming toward you
__2. structural formula N
H
H
B. H
__3. perspective drawing C. oxygen is an example
__4. ball-and-stick D. shows bonds going away from you
__5. space-filling
__6. closed arrows E.
__7. diatomic molecule AB. shows types and numbers of atoms only
__8. open arrows BC.
__9. molecular compound CD. compound composed of covalent bonds
__10. covalent bond N
DE. H H H
AC. created by sharing electrons
__11. The bond in a hydrogen molecule is a __.
A. single covalent bond B. double covalent bond
C. triple covalent bond D. ionic bond
__12. The bonds in a methane (CH4) molecule are all __.
A. single covalent bonds B. double covalent bonds
C. triple covalent bonds D. ionic bonds
__13. The bond in a nitrogen molecule is a __.
A. single covalent bond B. double covalent bond
C. triple covalent bond D. ionic bond
__14. The bonds between carbon and oxygen in carbon dioxide are __.
A. single covalent bonds B. double covalent bonds
C. triple covalent bonds D. ionic bonds
__15. A covalent bond in which one atom contributes both bonding electrons is called a ___ bond.
A. single covalent B. coordinate covalent
C. ionic D. polyatomic
__16. Which of the following is a property of most covalent compounds?
A. electrons transferred B. particles called formula units
C. low to no conductivity D. all are solids at room temperature
20
__17. What is another name for nonbonding pairs?
A. unshared pairs B. lone pairs
C. both A and B D. neither A nor B
__18. Why do atoms share electrons in covalent bonds?
A) to become ions and attract each other
B) to attain a noble-gas electron configuration
C) to become more polar
D) to increase their atomic numbers
__19. Once formed, how are coordinate covalent bonds different from other covalent bonds?
A) They are stronger. B) They are more ionic in character.
C) They are weaker. D) There is no difference.
Draw all dot structures for the following. (Each structure is worth 2 answers).
20. C2F2
FCCF
21. CS2
SCS
22. SiCl4
Cl
Cl C Cl
Cl
23. AsF3
F
F As F
21
8.4
Polar
Bonds
and
Molecules
Bond Polarity
• How do electronegativity values determine the charge distribution in a polar bond?
Covalent bonds differ in terms of how the bonded atoms __________ the electrons.
• The character of the molecule depends on the _________ and number of atoms
joined together.
• These features, in turn, determine the molecular properties.
The bonding pairs of electrons in covalent bonds are __________ between the nuclei
of the atoms __________ the electrons.
• When the atoms in the bond pull equally (as occurs when
_________________ atoms are bonded), the bonding electrons
are shared ___________, and each bond formed is a
covalent bond.
A covalent bond, is a covalent bond between atoms in
which the electrons are shared _______________.
- The ____________ electronegative atom attracts more strongly and gains a slight
______________ charge. The less electronegative atom has a slight positive charge.
Hydrogen has an electronegativity of ___, and chlorine has an
electronegativity of ____.
• (δ) is a lowercase delta symbol indicating a __________ charge.
• Cl atom, acquires a slight negative charge (δ-), and the H atom acquires a slight
positive charge (δ+).
δ+ δ– H—Cl
H—Cl
Describing Polar Covalent Bonds
These partial charges are shown as clouds of electron density.
This electron-cloud picture of hydrogen chloride shows that the _______________
atom attracts the electron cloud more than the hydrogen atom does.
The polar nature of the bond may also be represented by an arrow pointing to the
_________ electronegative atom.
The O—H bonds in a water molecule are also __________.
• The highly electronegative oxygen partially pulls the ___________ electrons away
from hydrogen.
22
• The ___________ acquires a slightly negative charge.
• The _____________ is left with a slightly positive charge.
The electronegativity difference (∆EN) between two atoms tells you what kind of
bond is likely to form.
Electronegativity Differences and Bond Types
∆EN type of bond Example
0.0 – 0.4
0.5 – 0.9 Nonpolar covalent H—H (2.20-2.20 = 0.0)
1.0 – 1.9 Polar covalent δ+ δ–
≥2.0
H—Cl (0.9)
Very polar covalent δ+ δ–
H—F (1.9)
Na+ Cl- (2.1)
Ionic
Describing Polar Covalent Molecules
The presence of a polar bond in a molecule often makes the _______ molecule polar.
• In order for the molecule to be polar, it must have _______________ and be
_______________________.
• This is what makes a molecule asymmetric: If a molecule has different
__________________ atoms AND/OR electrons pairs on the __________ atom.
• H2O is asymmetric and has polar bonds, therefore is a polar molecule
∆EN (O-H) = 3.44-2.20 = 1.24 (VPC)
since O has unshared pairs, it is a polar molecule
• CO2 is symmetric, so even though it has polar bonds, they “cancel” and it is a
nonpolar molecule
∆EN (C-O) = 3.44-2.55 = 0.89 (PC)
Since C has no unshared pairs and the molecule is symmetric, it
is a nonpolar molecule
23
8.4
Homework
(I
must
check
#s
11
and
12
before
you
scan)
Matching – Match the term/symbol on the right with the correct statement on the left. Some may be used
more than once and some may not be used at all.
__1. __ indicates a slight negative charge. A. asymmetric
__2. __ indicates a slight positive charge. B. intermolecular
__3. A __ is a bond where electrons are shared unequally. C. intramolecular
__4. In order for a molecule to be polar, it must have polar bonds and be __. D. ionic
__5. Polar bonds are represented by a __ or an arrow pointing toward the more E. lowercase delta
electronegative atom. symbol
__6. The attractive forces that hold particles together are call __ forces. AB. negative
__7. The less electronegative atom in a polar covalent bond gets a slight __ charge. AC. nonpolar covalent
__8. The more electronegative atom in a polar covalent bond gets a slight __ charge. AD. polar covalent
__9. The weaker attractive forces between particles are called __ forces. AE. positive
__10. When identical atoms bond and share electrons equally, this is called a __ BC. symmetric
bond.
BD. δ +
BE. δ-
Essay – Each is worth two answers.
11. What determines the degree of polarity in a bond? Distinguish between nonpolar covalent, polar
covalent, and ionic bonds in terms of relative polarity.
12. Explain this statement: Not every molecule with polar bonds is polar. Use CCl4 as the example.
24
Chapter 8 Practice Quiz
1-10 - Matching: Match the term with the description
__1. molecular formula A. indicate bonds coming toward you
__2. structural formula HN
B. H H
__3. perspective drawing C. oxygen is an example
__4. ball-and-stick D. shows bonds going away from you
__5. space-filling
__6. closed arrows E.
__7. diatomic molecule AB. shows types and numbers of atoms only
__8. open arrows BC.
__9. molecular compound CD. compound composed of covalent bonds
__10. covalent bond N
DE. H H H
AC. created by sharing electrons
11-25 – Multiple Choice: Choose the letter that best completes the statement or answers the question.
__11. The bond in a hydrogen molecule is a __.
A. single covalent bond B. double covalent bond C. triple covalent bond D. ionic bond
__12. The bonds in a methane (CH4) molecule are all __.
A. single covalent bonds B. double covalent bonds C. triple covalent bonds D. ionic bonds
__13. The bond in a nitrogen molecule is a __. C. triple covalent bond D. ionic bond
A. single covalent bond B. double covalent bond
__14. The bonds between carbon and oxygen in carbon dioxide are __.
A. single covalent bonds B. double covalent bonds C. triple covalent bonds D. ionic bonds
__15. A covalent bond in which one atom contributes both bonding electrons is called a ___ bond.
A. single covalent B. coordinate covalent C. ionic D. polyatomic
__16. Which of the following is a property of most covalent compounds?
A. electrons transferred B. particles called formula units
C. low to no conductivity D. all are solids at room temperature
__17. What is another name for nonbonding pairs?
A. unshared pairs B. lone pairs C. both A and B D. neither A nor B
__18. Why do atoms share electrons in covalent bonds?
A) to become ions and attract each other B) to attain a noble-gas electron configuration
C) to become more polar D) to increase their atomic numbers
__19. A __ symbol is used to indicate a partial charge.
A) ν B) δ
C) λ D) α
25
__20. When atoms in a covalent bond equally share electrons, they are considered __ bonds.
A) polar covalent B) nonpolar covalent
C) very polar covalent D) ionic
__21. In a polar covalent bond, the atom with a higher __ gets a slight negative charge.
A) atomic number B) electronegativity
C) atomic mass D) all of the above
__22. The polar nature of a bond can be represented by __.
A) a delta symbol B) an arrow pointing toward the more electronegative atom
C) differences in electron cloudsD) all of the above
__23. In a water molecule, the __ acquire a partial negative charge and the __ acquire a partial positive
charge.
A) hydrogens/oxygen B) oxygen/hydrogens
C) oxygen/oxygen D) hydrogens/hydrogens
__24. Which of the following covalent bonds is the most polar?
A) H-F B) H-C
C) H-H D) H-N
__25. In order for a molecule to be polar it must have __ and be __.
A) nonpolar bonds/symmetric B) nonpolar bonds/asymmetric
C) polar bonds/asymmetric D) none of the above
Draw all dot structures for the following.
26. OF2 27. CS2 28. SiCl4 29. PI3
26
Chapter
8
Study
Guides
8.1:
MOLECULAR
COMPOUNDS
–
Study
guide
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms that are introduced in this
section. Each blank can be completed with a term, short phrase, or number.
Every substance is either an element or a(n) ___1____. A 1.
compound is either ___2____ or ionic in nature. Most molecular 2.
compounds are composed of two or more ___3____. Molecules 3.
consisting of two atoms are ___4____ molecules. The chemical 4.
formula of a molecular compound is a ___5____. Molecular 5.
compounds tend to have ___6____ melting and boiling points, while 6.
ionic compounds tend to have ___7____ melting and boiling points. 7.
8.
A molecular formula shows how many ____8___ of each 9.
element a molecule contains, but it does not indicate the ___9____ of
the molecule.
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
10. A diatomic molecule contains two or three atoms.
11. Molecular compounds have relatively high boiling points.
12. The molecular structure of carbon dioxide is one carbon atom with two oxygen atoms on opposite sides of it.
13. Covalent bonds exist when combining atoms give up or accept electrons.
14. A molecule contains two atoms.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A Column B
15. molecule a. compound composed of molecules
16. molecular compound
17. covalent bond b. a molecule consisting of two atoms
18. diatomic molecule
19. molecular formula c. shows the kinds and numbers present in a molecule of a compound
d. joins atoms held together by sharing electrons
e. an electrically neutral group of atoms joined together by covalent
bonds
27
Part D Questions and Problems
Answer the following in the space provided.
20. A compound has a boiling point of 40°C. Is this compound most likely an ionic or a molecular compound?
21. Identify the number and kinds of atoms present in a molecule of each compound.
a. butane (C4H10)
b. fluorobenzene (C6H5F)
22. Classify each particle as an atom or a molecule.
a. CH4 d. He
b. Ne e. CO2
c. O2
28
8.2:
THE
NATURE
OF
COVALENT
BONDING
–
Study
guide
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms that are introduced in this section. Each blank
can be completed with a term, short phrase, or number.
When atoms share electrons to gain the ____1___ configuration of a noble gas, 1.
the bonds formed are ___2____. A ____3___ pair of valence electrons constitutes a 2.
____4___ covalent bond. Pairs of valence electrons that are not shared between 3.
atoms are called ___5____. Sometimes two or three pairs of electrons may be shared 4.
to give ____6___ covalent bonds. In some cases, only one of the atoms in a bond 5.
provides the pair of bonding electrons; this is a __7_____. ___8____ is required to 6.
break covalent bonds between atoms. The total energy required to break the bond 7.
between two covalently bonded atoms is known as the ___9____. 8.
9.
When it is possible to write two or more valid electron dot formulas for a 10.
molecule or ion, each formula is referred to as a ___1_0___.
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
11. The modern interpretation of resonance is that electron pairs rapidly flip back and forth between the
various electron dot structures.
12. The compound NH3 contains two double covalent bonds.
13. The chemical formulas of molecular compounds show the number and type of atoms in each molecule.
14. A molecule of bromine has six unshared pairs of electrons.
15. Carbon forms four single covalent bonds with other atoms.
16. A bond in which one atom contributes both bonding electrons is called a polyatomic covalent bond.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A Column B
17. single covalent bond a. a chemical formula that shows the arrangement of atoms in a
molecule or a polyatomic ion
18. structural formula
19. bond dissociation energy b. the amount of energy required to break a covalent bond between
20. polyatomic ion atoms
21. coordinate covalent bond
c. a tightly bound group of atoms that has a positive or negative
charge and behaves as a unit
d. a covalent bond in which one atom contributes both bonding
electrons
e. a chemical bond in which only one pair of electrons is shared by
two bonded atoms
29
8.4:
POLAR
BONDS
AND
MOLECULES
–
Study
guide
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms that are introduced in this section. Each blank
can be completed with a term, short phrase, or number.
When like atoms are joined by a covalent bond, the bonding electrons are 1.
shared ____1___, and the bond is _____2__. When the atoms in a bond are not the 2.
same, the bonding electrons are shared ___3____, and the bond is ____4___. The 3.
degree of polarity of a bond between any two atoms is determined by consulting a 4.
table of ___5____. The attractions between opposite poles of polar molecules are 5.
called ____6___. Another strong intermolecular attractive force is the ___7____, in 6.
which a hydrogen covalently bonded to a very ___8____ atom, such as ____9___, is 7.
also weakly bonded to an unshared electron pair of another electronegative atom. 8.
9.
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
10. In a polar covalent bond, the more electronegative atom has a slight positive charge.
11. In general, the electronegativity values of nonmetallic elements are greater than the electronegativity values of
metallic elements.
12. A molecule with polar bonds is dipolar.
13. Covalent compounds are network solids.
14. If the electronegativity difference between two atoms is greater than 2.0, they will form an ionic bond.
15. Dispersion forces are weaker than hydrogen bonds.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A Column B
16. nonpolar covalent bond a. a substance in which all of the atoms are covalently bonded to each other
17. polar covalent bond
18. polar molecule b. a bond formed when the atoms in a molecule are alike and the bonding electrons are share
19. van der Waals forces c. a term used to describe the weakest intermolecular attractions;
these include dispersion forces and dipole interactions
d. a bond formed when two different atoms are joined by a covalent bond and the bonding el
shared unequally
20. network solid e. a molecule in which one end is slightly positive and the other end is
slightly negative
30
Part D Questions and Problems
Answer the following in the space provided.
21. Arrange the following intermolecular attractions in order of increasing strength: dipole interactions, dispersion forces, and
hydrogen bonds.
22. State whether the following compounds contain polar covalent bonds, non-polar covalent bonds, or ionic bonds, based on their
electronegativities.
a. KF a.
b. SO2 b.
c. NO2 c.
d. Cl2 d.
31
Chapter
10
-‐ Chemical
Quantities
10.1 – The Mole: A Measurement of Matter
Objectives
• Relate Avogadro’s number to a mole of a substance
• Calculate the mass of a mole of any substance
• Describe methods of measuring the amount of something
• Compare and contrast the atomic mass of an element and its molar mass
Vocabulary • representative particle
• molar mass
• mole (mol)
• Avogadro’s number
10.2 – Mole-Mass and Mole-Volume Relationships
Objectives
• Convert the mass of a substance to the # of moles of a substance, and the number of moles of a substance to mass
• Calculate the volume of a quantity of gas at STP
Vocabulary
• Avogadro’s hypothesis
• standard temperature and pressure (STP)
• molar volume
Key Equations
• mass (grams) = number of moles × mass (grams)
1 mole
• moles = mass (grams) × 1 mole
mass (grams)
• grams = grams × 22.4 L
mole L 1 mole
• volume of gas = moles of gas × 22.4 L
1 mole
10.3 – Percent Composition and Chemical Formulas
Objectives
• Calculate the percent by mass of an element in a compound
• Interpret an empirical formula
• Compare and contrast empirical and molecular formulas
Vocabulary
• percent composition
• empirical formula
Key Equation
% mass of element = mass of element × 100%
mass of compound
32
10.1
Chemical
Measurements
Atomic mass - the mass of an atom in atomic mass units (_______) = 1/12 the mass of 1 Carbon-12 atom.
(the decimal # on the periodic table) ** We will round atomic masses to 1 decimal place**
Ex. Cd = 112.4 amu, Bi = 209.0 amu
formula mass - sum of masses of all atoms in a compound.
Ex. CCl4 1 atom C ✕12.0 amu/atom C = ______ amu
+ 4 atoms Cl ✕ 35.5 amu/atom Cl = _______ amu
molecular mass of CCl4 = _______ amu
Representative Particle (RP) - atoms, ion, or molecules.
- most elements exist as ____________ (Fe, Au, Na, Hg)
- seven elements exist as _________________ molecules (H2,N2,O2,F2,Cl2,Br2,I2)
- ____________________ compounds exist as molecules (H2O,CO2)
- ionic compounds exist as __________________________________ (NaCl, CuCl2)
A ___________ (mol) of any element is the # of atoms of that element equal to the # of atoms in exactly
12.0 g of carbon-12. 1 mol = ___________________________ RP (called Avogadro’s Number). - the mole
shows the relationship between the atomic mass unit and ________________.
- the mass in grams of 1 mole of a substance is _______________________ equal to its atomic mass or
formula mass in amu.
- _______________________ is the mass (in grams) of 1 mole of a substance
Mole relationships
How many atoms are in a molecule of ammonia (NH3)?
__ total atoms– three __ and one __
1 mole of NH3 has 1 mole of molecules, how many moles of atoms are in a mole of ammonia?
Since there are 4 atoms in 1 molecule, there are 4 moles of atoms in 1 mole of ammonia (3 mol H, and 1 mol N)
(this relationship would work for ionic compounds broken into its ions or atoms that make up the ions)
33
10-‐1
Worksheet
–
What
is
a
mole?
Exponential Numbers
To enter very large or very small numbers on a scientific calculator, use the exponential key. The exponential key on
your calculator may be labeled [EE], [EXP], or [EEX]. In this worksheet, the exponential key will be represented by
[EE].
Suppose you want to enter the exponential number 2.94 × 1024 on your calculator. First enter the coefficient
[2][.][9][4]. Then press [EE] and enter the exponent [2][4]. The display reads 2.9424. Notice that the “×10” part of
the exponential form is not entered on (or displayed by) the calculator. Pressing the exponential key indicates to the
calculator that the next number entered is a power of 10.
If a number written in exponential form has a negative exponent, use the sign-change key [+/−] (not the
subtraction key!) to make the sign of the exponent negative.
For example, to enter the number 3.08 × 10−12 on your calculator, first enter the coefficient [3][.][0][8]. Then
press [EE] followed by [+/−]. Notice how the display changes from 3.0800 to 3.08−00. Finally, enter [1][2], the value
of the exponent. The display now reads 3.08−12.
You can practice entering exponential numbers on a calculator by trying the following examples. The keystrokes
are listed after each exponential number. The numbers displayed by your calculator should match those in the last
column.
Numeric Entry Keys Display
a. 6.02 × 1023 [6] [.] [0] [2] [EE] [2] [3] 6.0223 Avogadro’s Number
b. 1.11 × 10−5 [1] [.] [1] [1] [EE] [+/−] [5] 1.11−05
What is a mole?
A mole is equal to 6.02 × 1023 representative particles (RP) of a substance. From this equality we can get two
!.!" ×! ! !""!#" !".
conversion factors: ! !"# and We can use either of these conversion factors to convert
!.!" × !"!" !"
from either moles to RP or vice versa.
Converting Representative Particles to Moles
Ex. Magnesium is a light metal used in the manufacturing of aircraft, automobile wheels, and tools. How many
moles of Mg is 1.25 × 1023 atoms Mg?
W – mol Mg
H –1.25 × 1023 atoms Mg
1) How many moles is 2.80 × 1023 atoms of silicon?
W – mol Si
H - 2.80 × 1023 atoms Si
2) How many moles is 2.17 × 1023 molecules of Br2?
W – mol Br2
H - 2.17 × 1023 molecules Br2
34
Converting Moles to Number of Atoms (in a Compound)
Ex. Propane (C3H8) is a gas used for cooking and heating. How many total atoms are in 2.12 mol of C3H8?
W – total atoms (C + H)
H – 2.12 mol C3H8
2.12 mol C3H8 ×
!.!!" !× "!#" !!"" !"
×
!! !"!# !$"% ! (!!!!!!!!)
=
1.40 × 1025 atoms total
3) How many atoms are in 1.14 mol of SO3?
4) How many C atoms are in 2.12 mol of C3H8? How many H atoms are in 2.12 mol of C3H8?
Finding Molar Mass of a Compound
Ex. The decomposition of hydrogen peroxide (H2O2) provides sufficient energy to launch a rocket. What is the
molar mass of H2O2?
W – M (molar mass) of H2O2
H – 1.0 g H/1 mol H
H – 16.0 g O/1 mol O
5) Find the molar mass of phosphorus trichloride.
6) What is the mass of 1.00 mol of sodium hydrogen carbonate?
35
10-‐2
Mole
Conversions
Mass and Moles
- if you know the mass of a substance, you can calculate the # of moles (and vice versa).
Remember this:
1: to convert between moles and mass, you will need to calculate molar mass.
2: Set up the problem placing what you have (what is given in the problem) and convert to what you want.
Ex. How many grams are in 7.50 moles of calcium hydroxide? 1: formula - Ca(OH)2
molar mass =
element # of atoms × mass (g) total mass (g)
Ca 1 × 40.1
O 2 × 16.0
H 2 × 1.0
Ca(OH)2 molar mass =
2:
3. Find the molar mass of carbon monoxide (CO).
4. Stomach acid is made up of hydrochloric acid (HCl). What is the molar mass of HCl?
5. Find the mass of 0.650 mol P2O5.
6. A bottle of NaNO3 contains 100.0 g of the compound. How many moles of NaNO3 does it contain?
- this conversion is simpler, because you do not have to calculate molar mass. Just use Avogadro’s # (6.02×1023
RP) (See Fig. 10-18). See sample problem #4 and complete practice problems 7 and 8 on p. 327
36
7. Determine the number of atoms in 0.36 mol Al.
8. How many moles of sodium carbonate (Na2CO3) contain 7.9×1024 formula units (u)?
Multi-step Conversions
Suppose you are given 4.8 g of Ca, how would you calculate the # of Ca atoms?
- Do it step-by-step.
1) convert to moles
2) convert from moles to atoms
This can be done in one step:
9. How many formula units of NaHCO3 are in 1.8 g of sodium bicarbonate (baking soda)?
10. If you burned 4.0×1024 molecules of natural gas, or methane (CH4), during a laboratory experiment, what
mass of methane did you burn?
Moles and Gases
Volume of a mole of gas (see Fig. 10-22, p. 330) At STP (standard temp. {0℃} and pressure {1 atm}),
one mole of any gas occupies 22.4L. 22.4 L is called the molar volume of a gas. 22.4 L is also equal to
6.02×1023 RP of that gas.
Example: What is the vol. of 5.40 mol O2?
37
11. A room with a volume of 4000. L contains how many moles of air at STP?
12. A chemical reaction produces 0.82 moles of oxygen gas. What volume will that gas occupy at STP?
10.1/10.2
Homework
(you
must
show
work
for
credit)
__1. Calculate the molar mass of calcium fluoride (CaF2). Show your work.
A. 118.2 g/mol
B. 99.2 g/mol
C. 78.1 g/mol
D. 59.1 g/mol
E. 50.0 g/mol
__2. Calculate the molar mass of potassium permanganate, KMnO4. Show your work.
A. 52.0 g/mol
B. 70.0 g/mol
C. 110.0 g/mol
D. 158.0 g/mol
E. 176.0 g/mol
__3. The mass of 1.63 × 1021 Si atoms is __. Show your work.
A. 2.71 × 10-23 g
B. 4.58 × 1022 g
C. 28.1 g
D. 1.04 × 104 g
E. 0.0761 g
38
__4. Calculate the number of moles in 17.8 g of the antacid magnesium hydroxide, Mg(OH)2. Show your work.
A. 3.28 mol
B. 2.32 mol
C. 0.431 mol
D. 0.305 mol
E. 0.200 mol
__5. Calculate the mass in grams of 3.65 × 1020 molecules of sulfur trioxide (SO3). Show your work.
A. 0.000606 g
B. 0.0291 g
C. 0.0486 g
D. 20.6 g
E. 1650 g
__6. What is the mass, in grams, of one copper (Cu) atom? Show your work.
A. 1.06 × 10-22 g
B. 63.6 g
C. 1 amu
D. 1.66 × 10-24 g
E. 9.48 × 1021 g
__7. One mole of iron __.
A. is heavier than one mole of lead (Pb).
B. is 77.0 g of iron.
C. is 26.0 g of iron.
D. weighs the same as one mole of lead.
E. is none of these.
__8. Which of these quantities does not represent 1.00 mol of the indicated substance? Explain your reasoning.
A. 6.02 × 1023 C atoms
B. 26.0 g Fe
C. 12.0 g C
D. 65.4 g Zn
E. 6.02 × 1023 Fe atoms
__9. How many silicon atoms are there in 1.00 g of silicon? Show your work.
A. 1 atom
B. 0.0356 atoms
C. 2.57 × 1023 atoms
D. 2.14 × 1022 atoms
E. 1.75 × 1025 atoms
__10. Calculate the number of moles of xenon in 12.0 g of xenon. Show your work.
A. 1.00 mol
B. 0.0457 mol
C. 0.183 mol
D. 7.62 × 10-3 mol
E. 0.0914 mol
39
10-‐3
Empirical
and
Molecular
Formulas
Percentage Composition- the mass of each element in a compound compared to the entire mass of the
compound and multiplied by 100%
Calculating % Composition (Using the given formula)
1- calculate the molar mass of the compound
2- calculate the mass of each element in the compound
3- (mass of each element/mass of compound) 100% Ex: H2O
- Using experimental data
1- measure the mass of the sample
2- the sample is chemically separated or decomposed
3- the masses of the components are determined
4- percentage composition calculated using this data
(mass of each element/mass of sample)×100%
(see practice problem 13 for example)
13. Find the % comp. of a compound that contains 2.30 g of Na, 1.60 g of O, and 0.100 g of H in a 4.00 g
sample of the compound.
14. A sample of unknown compound with a mass of 0.562 g has the following % comp.: 13.0% C, 2.20% H,
and 84.8% F. When the compound is decomposed into its elements, what mass of each element would be
recovered?
40
Determining Empirical Formula
- it is the “opposite” of percentage composition
- an empirical formula gives the simplest whole-number ratio of the atoms of the elements in a compound
Determining Empirical Formulas
1- assume 100 g of the compound (this makes the mass in grams of each element equal in # to the percentage)
2- the ratio of atoms in formula is the ratio of moles in a mole of the compound
2a- subscripts indicate the # of moles of atoms in a compound
3-convert mass of each element to # of moles of atoms
4- divide ALL moles of elements by the smallest mole value
5- round all to nearest whole number
(if .1 or .9, round to nearest # – if any other decimal, multiply ALL by that # to get whole #s)
15. Determine the empirical formula of a compound containing 2.128 g Cl and 1.203 g Ca.
16. Determine the empirical formula of a compound containing 7.30 g Na, 5.08 g S, and 7.62 g O.
10.3
Homework
This HW will be checked by hand, no GradeCam!
1. Find the percent composition of a compound containing tin and chlorine if 18.35 g of the compound contains 5.74 g of tin.
2. If 3.907 g of carbon combines completely with 0.874 g of hydrogen to form a compound, what is the percent composition of this
compound?
3. How many grams of aluminum are in 25.0 g of aluminum oxide (Al2O3)?
4. Determine the empirical formula from the percent composition: 10.0% C, 0.80% H, 89.1% Cl
41
Chapter
10
Practice
Quiz
a. the mass in amu of an atom of an element
b. the formula that gives the lowest whole-number ratio of the
1-9 – Matching
__1. atomic mass elements in a compound
__2. Avogadro’s number c. an atom, a formula unit, or a molecule
d. 6.02 × 1023 representative particles of a substance
__3. empirical formula e. the volume occupied by a mole of any gas at STP (22.4 L)
__4. molar mass
__5. molar volume f. the percent by mass of each element in a compound
__6. mole (mol) g. the number of representative particles in a mole of a substance
__7. percent composition
__8. representative particle h. the mass of a mole of a substance
__9. standard temperature and pressure i. 0°C and 1 atmosphere
Problems –Each will count as three answers. YOU MUST SHOW WORK FOR CREDIT.
1 mole = 6.02 × 1023 RP = 22.4L at STP = molar mass
10. Calculate the molar mass of CaSO3.
W – molar mass
H – 1 mol Ca
H – 1 mol S
H – 3 mol O
11. How many moles are in 140 g Al?
W – mol Al
H – 140 g Al
12. How many liters are in 5.83 mol of O2 gas at STP?
W – L O2
H – 5.83 mol O2
More on the back!
42
13. How many moles are in is 2.68 × 1023 formula units of AlF3?
W - mol AlF3
H – 2.68 × 1023 u AlF3
14. Find the mass of 143L of C2H4 at STP.
W – g C2H4
H – 143 L of C2H4
15. What is the empirical formula of a compound with a composition by mass of 35% P and 65%F?
W – PxFy
H – 35% P
H – 65% F
43
Chapter
10
Study
Guides
10.1:
The
Mole:
A
Measurement
of
Matter
–
Study
guide
Part A Completion
Use this completion exercise to check your knowledge of the terms and your understanding of the concepts introduced in this
section. Each blank can be completed with a term, short phrase, or number.
Chemists relate units of counting, of mass, and of 1.
volume to a single quantity called the ____1___. The number of 2.
representative particles in a mole of a substance is _____2__. 3.
4.
To find the mass of a mole of a compound, scientists add 5.
together the ___3____ of the atoms making up the compound.
When you substitute the unit grams for amu, you obtain the
____4___ of the compound. There are ___5____ representative
particles in a mole of any substance.
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
6. A mole of a pure substance contains 6.02 × 1023 atoms.
7. The representative particle of a compound is the molecule.
8. A mole of CCl4 is composed of one atom of carbon and four atoms of chlorine.
9. A mole of carbon atoms has a mass approximately three times as great as the mass of a mole of helium atoms.
10. The molar mass of nitrogen gas is 14.0 g.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A Column B
11. Avogadro’s number a. the atoms, molecules, or ions present in a substance
12. molar mass b. 6.02 × 1023
13. mole c. the mass of one mole of a substance
14. representative particles d. SI unit that measures the amount of a substance
44
Part D Problems
Solve the following problems in the space provided. Show your work.
15. How many moles of Pb is 9.3 × 1015 atoms of Pb?
16. What is the molar mass of ethane, C2H6?
17. Find the mass of 3.65 × 10-2 mol K2SO4.
18. How many representative particles are in 2.5 mol H2O2?
45
10.2:
Mole-‐Mass
and
Mole-‐Volume
Relationships
–
Study
guide
Key Equations
• mass (grams) = number of moles × mass (grams)
1 mole
• moles = mass (grams) × 1 mole
mass (grams)
• grams = grams × 22.4 L
mole L 1 mole
• volume of gas = moles of gas × 22.4 L
1 mole
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
6. One mole of any gas occupies a volume of 22.4 L.
7. For a substance of known molar mass, the number of moles of a sample can be calculated from the mass of the sample.
8. The volume occupied by one mole of a gas is dependent on the molar mass of the gas.
9. The volume of a gas at STP can be calculated from the number of molecules of the gas.
46
Part C Matching Column B
Match each description in Column B to the correct term in Column A.
Column A
10. molar mass a. 22.4 L of a gas at STP
11. standard temperature b. 101.3 kPa or 1 atm
12. molar volume c. 0°C
13. standard pressure d. mass (in grams) of one mole of a substance
14. molar road map
e. a means of relating mass, number of representative
particles, and gaseous volume of a substance
Part D Problems
Solve the following problems in the space provided. Show your work.
15. What is the density of N2O, a gas, at STP?
16. What is the mass of two moles of NaCl?
17. How many moles are in 16 grams of O2?
18. What is the volume of 16 grams of O2 at STP?
47
10.3:
Percent
Composition
and
Chemical
Formulas
–
Study
guide
Key Equation
• % mass of element = mass of element × 100%
mass of compound
Part A Completion
Use this completion exercise to check your knowledge of the terms and your understanding of the concepts introduced in this
section. Each blank can be completed with a term, short phrase, or number.
The ___1____ of a compound is the percent by mass of each element in a 1.
compound. The percent by mass of an element in a compound is the number of 2.
grams of the element per ____2___ g of the compound, multiplied by 100%. To 3.
calculate the percent by mass of an element in a known compound, divide the mass 4.
of the element in one mole by the ____3___ and multiply by 100%. 5.
6.
A(n) ___4____ formula represents the lowest ___5____ ratio of the elements in
a compound. It can be calculated from a compound’s percent composition. The
____6___ formula of a compound is either the same as its empirical formula, or it is
some whole-number multiple of it.
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
7. It is necessary to know the formula of a compound in order to calculate its percent composition.
8. If the percent by mass of carbon in methane, CH4, is 75%, then 100 grams of methane contain 25.0 grams
of hydrogen.
9. The formula for methane, CH4, is both a molecular and an empirical formula.
10. The empirical formula for glucose, C6H12O6, is C2H4O2.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A Column B
11. percent composition a. describes the actual number of atoms of each element in
12. empirical formula a molecule of a compound
13. molecular formula
b. the lowest whole-number ratio of atoms of the elements
in a compound
c. the percent by mass of each element in a compound
48
Part D Problems
Solve the following problems in the space provided. Show your work.
14. What is the percent composition of each of the following?
a. Cr2O3 c. HgS
b. Mn2P2O7 d. Ca(NO3)2
15. Determine the empirical formula of the compound with the percent
composition of 29.1% Na, 40.5% S, and 30.4% O.
16. How many kilograms of iron can be recovered from 639 kilograms of the ore Fe2O3?
49
Ionic
or
Covalent
Bonding
Lab
Purpose: Some properties may be useful to predict the type of bonding in a substance. These properties are
phase at room temperature, melting point, solubility in water, and electrical conductivity. In this experiment
you will find how these properties vary in ionic and covalently bonded substances.
Procedure: Part One
1.) Place a tiny amount of each sample (<1g) of each substance in each well of the dish provided. Put one
sample in one well only. Do not put in the distilled water.
2.) In the data table record the phase at room temperature.
3.) Test each substance for electrical conductivity using the meter provided. Record your results in the data
table. DO NOT add water until the dry substances are tested first.
4.) Place several drops of distilled water into the wells with a solid sample. Do not put water into the liquid
samples. Mix the samples so that each has a chance to dissolve in the water.
5.) Determine which substance is soluble. HINT: which substances dissolved in the water and which did not?
Record this in the data table under “solubility”. The substance dissolves if the original substance is no longer
visible. The water solution may have a color but should be transparent.
6.) Test each substance a second time using the conductivity meter provided. Record which substance will
conduct electricity in the data table.
Part Two: Testing the Melting point of the solid substances.
1.) Make a 2 small foil cups by wrapping a piece of aluminum foil around your thumb. Your cups should be
about the size of the end of your thumb.
2.) Put a small amount (1 gram) of one solid into a small foil cup.
3.) Heat each cup over the Bunsen burners for about 30 seconds. Record in the data table which substance
melts, decomposes or in which nothing happens at all.
Substance Ionic or Solid or Melting Point Electrical Conductivity Solubility
Covalent liquid (high or low) Conductivity with water (yes or no)
Distilled
water Bond (Yes or no) (yes or no)
NaCl
KCl
Sugar
Oil
Ethanol
Starch
Glycerine
CaCl2
50
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Bonding and Moles unit
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