Unit 3: The Elements
Ionic and Covalent Compounds
Bonding
(Ch 7, 8, & 9)
October 3 – October 21
Science Interactive Notebook
Grant Union High School
2011-2012
Science GEMS of Wisdom
Writing – a way to clarify your thoughts
Collaboration – a way to share your thoughts
about
Inquiry – a way to expand your thoughts
Reading – a way to gain new knowledge to think
GEMS: Scientists & Discoveries, Types of Problems to Solve, Formulas & Constants, Specialized Vocabulary,
Major ideas & Facts and Discoveries, Types of Problems to Solve, Formulas & Constants, Specialized
Vocabulary.
Science GEMS of Wisdom
Writing – a way to clarify your thoughts
Collaboration – a way to share your thoughts
about
Inquiry – a way to expand your thoughts
Reading – a way to gain new knowledge to think
GEMS: Scientists & Discoveries, Types of Problems to Solve, Formulas & Constants, Specialized Vocabulary,
Major ideas & Facts and Discoveries, Types of Problems to Solve, Formulas & Constants, Specialized
Vocabulary
Science GEMS of Wisdom
Writing – a way to clarify your thoughts
Collaboration – a way to share your thoughts
about
Inquiry – a way to expand your thoughts
Reading – a way to gain new knowledge to think
GEMS: Scientists & Discoveries, Types of Problems to Solve, Formulas & Constants, Specialized Vocabulary,
Major ideas & Facts and Discoveries, Types of Problems to Solve, Formulas & Constants, Specialized
Vocabulary
Interactive Notebook Table of Contents
Pg Date Left Side Items
CK
Pg
Date Right Side Items
0
GEMS of Wisdom
1
Table of Contents
2
Score Sheet
3
Scoring Rubrics/Safety
Contract
4
5
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8
9
10
11
12
13
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41
CK
Interactive Notebook Score Sheet
Unit:
Quizzes/Formatives Date
Name of Special Assignment
Quarter
Score/Max
Score
Name
Retake Needed
(yes or no)
Date Scored Score/Max
Peer Initial
Peer Initials
Parent Initial
Teacher Initials
Five Point Scoring Rubric
5 Points—(a WOW product)
• all of the requirements are evident and EXCEEDED
• the product is VERY neatly done and EXTREMELY well organized
• the product shows LOTS of creativity and is colorfully illustrated
• completed on time
4 Points—(What is EXPECTED)
• all the requirements are evident
• the product is neatly done and well organized
• the product shows creativity and is colorfully illustrated
• completed on time
3 Points—(Almost What is EXPECTED)
• the requirements are evident (maybe 1 or 2 are missing)
• the product is neatly done and organized
• the product shows some creativity and is illustrated
• completed on time
2 Points—(Sort of What is EXPECTED)
• the requirements are evident (maybe 3 or 4 are missing)
• the product is done and sort of organized
• the product shows little creativity and is illustrated
• completed on time
1 Points—(Two or More parts is missing)
• MANY of the requirements are NOT PRESENT
• The product is VERY POORLY done and POORLY organized
• The product shows little TO NO creativity and THE illustrations IS POORLY DONE
0 Points—(Does not meet Standards)
• Unscorable or no product
I understand that I am responsible for my
personal safety, as well as others in the class. I
will follow the safety rules below and
additional instructions given by my teacher. I
understand that the following behaviors will
result in a “time out” from lab, a referral, a parent
conference, detention, or removal from the class:
 Failure to abide by class safety rules
 Behavior that causes injury to myself or another person
 Intentional damage or theft of classroom materials or
equipment
I will behave responsibly at all times.

Understand and follow written and verbal instructions

Avoid eating, drinking or chewing gum in the classroom
or lab at all times

Avoid any actions (horseplay, running) that might injure
myself or another student

Perform only authorized experiments
I will work safely at all times.

Wear closed toed shoes and tie my hair back in lab

Understand emergency procedures

Know location and use of safety equipment

Wear safety goggles or other protective equipment when
instructed to do so

Dispose of chemicals and broken glass properly

Wash hands after working with hazardous materials

Report accidents or spills to the teacher immediately
I will be cautious when working with laboratory equipment.

Understand the proper use and care of equipment

Use equipment only for its purpose as specified Treat all
lab equipment – carefully
I will be especially aware and cautious when dealing with
heat or flames.

Follow specific directions when using matches, candles, a
hotplate, or burner

Know the proper use of hot plates and Bunsen burners
before using them

Not leave hot objects unattended

Keep flammable materials (alcohol, hairspray) away from
flames
I will be a good citizen in the classroom.

Clean my work area and return supplies & equipment

Dispose of all chemicals/specimens according to
instructions

Not touch equipment or supplies until instructed

Stay out of off-limits areas

Take responsibility for broken glassware, equipment, or
spills
Student Signature _______________________________
Parent’s Review __________________________________
California Standard Chemical Bonds
2.
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. As a basis for understanding this
concept:
a.
Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by
exchanging electrons to form ionic bonds.
b.
Students know chemical bonds between atoms in molecules such as H2 , CH4 , NH3 , H2 CCH2 , N2 , Cl2 , and many
large biological molecules are covalent bonds.
c.
Students know salt crystals, such as NaCl, are repeating patterns of positive and negative ions held together by
electrostatic attraction between ions.
d.
Students know the atoms and molecules in liquids move in a random pattern relative to one another because
the intermolecular forces are too weak to hold the atoms or molecules in a solid form.
e.
Students know how to draw Lewis dot structures.
f.
* Students know how to predict the shape of simple molecules and their polarity from
Lewis
dot structures.
g.
* Students know how electronegativity and ionization energy relate to bond
formation.
h.
* Students know how to identify solids and liquids held together by van der Waals
forces or hydrogen bonding and relate these forces to volatility and boiling/ melting
point temperatures.
Think Pair Share – What are some meanings for the term bonding
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Metallic
Ionic
Covalent
Bonding Introduction (215-275 textbook)
Bonding occurs in order to become more stable. When a bond is formed energy is released because
the resulting compound is more stable and at a lower energy state.
When energy is released the term exothermic is used.
Exo means______________.
When bonds are broken energy is absorbed and the term endothermic is used. Endo means._________
A chemical bond results from the simultaneous attraction of electrons by two nuclei. Bonding only
involves electrons. Atoms bond together to form OCTETS (eight valence electrons) the most stable
state. The exceptions are hydrogen and helium which need only two electrons to fill their outer orbitals
(shells)
Standards Review Chemistry 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 depending on the
atom’s electronegativity. Electronegativity - the affinity for electrons. Highest: Fluorine 4.0.
A covalent bond between non-metals appears as a shared pair of electrons contained in a region of
overlap between two atomic orbitals, in the localized electron model. Atoms (usually nonmetals) of similar
electronegativities can form covalent bonds to become molecules. In a covalent bond, therefore, bonding
electron pairs are localized in the region between the bonded atoms.
In metallic bonds between metals valence electrons are not localized to individual atoms but are free to
move to temporarily occupy vacant orbitals on adjacent metal atoms. For this reason metals conduct
electricity well. This kind of bonding is also called the sea of electrons.
Ionic bonding between metals and non-metals; occurs when an electron from an atom with low
electronegativity (e.g., a metal) is removed by another atom with high electronegativity (e.g., a
nonmetal), the two atoms become oppositely charged ions that attract each other.

Metals tend to lose electrons and form positive ions called cations. Ions
formed are smaller than the neutral atoms because they have lost
electrons: Ionic radii < than atomic radii.

Nonmetal tend to gain electrons and from negative ions, called anions. Ions
formed are larger than the neutral atoms because anions gain electrons:
Ionic radii > atomic radii.
Chemical bonds between atoms can be almost entirely covalent (nonpolar), almost entirely ionic or in
between these two extremes. The triple bond in nitrogen molecules (N2) is nearly 100 percent covalent
as are other diatomic molecules. A salt such as sodium chloride (NaCl) has bonds that are nearly
completely ionic. However, the electrons in gaseous hydrogen chloride are shared somewhat unevenly
between the two atoms. This kind of bond is called polar covalent.

Ionic bonds - formed between metal and nonmetal; created by a transfer of electrons;
electronegativity difference > 1.7

Covalent bond - formed by the sharing of electrons between nonmetals; electronegativity
difference < 1.7. Nonpolar covalent bonds have electronegativity differences <0.3 and Polar
covalent bonds have electronegativity >0.3 and <1.7

Diatomic molecules are considered to have NONPOLAR covalent bonding. i.e. N2

Exception to 1.7 rule: METAL hydrides are ionic! ex. NaH
Electrostatic Forces and Crystalline Structure
Order the elements by increasing
electronegativity
Na, Rb, O, F
Order the elements by decreasing
electronegativity
Br, I, F, Sr, Ba
Electron affinity and Electronegativity
Elements with a high electron affinity also have high electronegativities. One term, electron affinity,
relates to attraction to gain an electron this is affinity. The other term electronegativity refers to the
pull of electrons once a bond has formed.
California Standard 2c. Students know salt crystals, such as NaCl, are repeating patterns of
positive and negative ions held together by electrostatic attraction
Description
The energy that holds ionic compounds together, called lattice energy is caused by the electrostatic
attraction of cations, which are positive ions, with anions, which are negative ions. To minimize their
energy state, the ions form repeating patterns that reduce the distance between positive and negative
ions and maximize the distance between ions of like charges.
Prepare a labeled diagram to illustrate the description and how water dissolves ionic crystal lattice.
Lewis Dot Diagrams and Bonding
Draw the Lewis Dot Diagram below the following elements:
1. Magnesium
2. Potassium
3. Nitrogen
4. Neon
5. Aluminum
6. Sulfur
7. Lithium
8. Argon
9. Helium
10. What is an ionic bond?
11. What kinds of elements do ionic bonds form between?
12. Would Potassium (K) form an ionic bond with Fluorine (F)? Why or Why not?
13. Would Calcium (Ca) form an ionic bond with Neon (Ne)? Explain why or why not.
14. Would Nitrogen (N) form an ionic bond with Oxygen (O)? Why or Why not?
15. Would Calcium become an anion (-) or a cation (+)? Explain why.
16. Would Phosphorus (P) become an anion (-) or a cation (+)? Explain why.
2e. Students know how to draw Lewis dot structures
A Lewis dot structure shows how valence electrons and covalent bonds are arranged
between atoms in a molecule. Knowledge of the periodic tables allows the determination of
the number of valence electrons for each element in Groups 1 through 3 and 13 through 18.
Carbon, for example, would have four valence electrons. Lewis dot diagrams represent each
electron as a dot or an x placed around the symbol for carbon, which is C. A covalent bond is
shown as a pair of dots, or x’s, representing a pair of electrons or a line. For example, a
Lewis dot diagram for methane, which is CH4, would appear as shown in Figure 3.
DIRECTIONS: For each of the following, draw the electron dot structure and identify the charge
(oxidation State) of the ion it will form.
1.
Ba
2.
O
3.
Ga
4.
Cl
5.
He
6.
P
7.
Tl
8.
Li
9.
Al
10.
K
11.
Ca
12.
Ar
SECTION 2. Predict the chemical formulas for each of the following ionic compounds.
Example: Aluminum reacts with oxygen in air to form a whitish coating. Predict the formula for
the aluminum oxide formed.
13. K and O
15. Hg+2 and P
14. Al and S
SECTION 3. CIRCLE THE ELEMENT IN EACH OF THE FOLLOWING IONIC COMPOUNDS WITH THE HIGHEST
ELECTRONEGATIVITY.
16.
MgCl2
17.
AlN
18.
SO4-2
SECTION 4. CIRCLE WHICH ELEMENT IN THE FOLLOWING PAIR THAT IS MOST LIKELY TO LOSE AN
ELECTRON BECAUSE IT HAS A LOWER IONIZATION ENERGY.
19.
As or Se
20.
Kr or Ne
21.
Li or F
22.
H or He
23.
Fr or Rb
24.
B or O
25.
C or Ge
26.
Br or Sc
27.
Be or Mg
NAMING – IONIC COMPOUNDS. BINARY TWO ELEMENTS, TERNARY MORE THAN TWO ELEMENTS
Compound Name
Classification
Chemical Formula
Fe(OH)3
Pb3N4
Al2(SO4)3
Hg(OH)2
COMPOUNDS CRISS-CROSS METHOD
Write the correct formula of the ionic compounds formed between the positive and negative ions.
Cl
K
Ca
Al
Na
Cu+1
Fe+3
NH4
O
N
P
KCl
K3P
Potassium
Chloride
Potassium
Phosphide
S
Write formula and name the ionic compounds formed
CO3
K
Ca
Al
Na
Cu+1
Fe3+
NH4
NO3
SO4
PO4
OH
Ionic Bonding Review
1. Key Idea/s:
1d. Students know how to
use the periodic table to
determine the number of
electrons available for
bonding.
b. Students can determine
the number of valence
electrons present in a
given atom based on the
element’s location in the
Periodic Table.
How do we know how many valence
electrons are in elements
2. Key Idea/s:
Students can explain that
only the outermost
electrons [or valence
electrons] are involved in
bonding.
What are valence electrons?
3. Key Idea/s:
Students know that either
positive or negative ions
form and what these ions
are called
How does an atom become a positive
ion?
How does an atom become a negative
ion?
How do ions form neutral compounds?
When do ionic compounds form
How many valence electrons do elements
desire?
Examples:
2a has________
3a has________
4a has________
5a has________
6a has________
Which elements do not follow this
trend?
What kind of elements form ionic
bonds?
The number of valence electrons are used
to determine what?
4. Key Idea/s:
Ionic Radii are different
from Atomic radii
Explain why cations get
smaller than their atomic
form and anions get
bigger than their atomic
form
5. Key Idea/s:
2a. Students know atoms
combine to form molecules
by exchanging electrons to
form ionic bonds.
6. Key Idea/s:
2c. Students know salt
crystals, such as NaCl, are
repeating patterns of
positive and negative ions
held together by
electrostatic attraction.
Discuss how dissolution breaks ionic bonds.
7. Key Idea/s:
Students know the
characteristics of ionic
compounds
What characteristics do ionic compounds share?
8. Key Idea/s
Students can explain how
electrical conductivity
occurs
How do ionic compounds conduct
electricity?
9. Key Idea/s:
Students can determine
the bonding capacities of
each atom[or element] by
examining the combining
ratios of the elements in
the compounds it forms
How many potassium ions are needed to
balance the charge of one sulfide ion?
10. Key Idea/s:
What are polyatomic ions?
Predict the formulas for calcium
chloride and potassium phosphate?
Examples:
11. Key Idea
2e. Students know how
to draw Lewis
structures
Draw KCl
Draw MgCl2
12. Key idea
Students know how to
name ionic compounds
Name MgO
Na2SO4
MElizabeth/HS_Chem_Ionic_Bond_Review/2010
Standards Review Chemistry 2b.
Students know chemical bonds between atoms in molecules
such as H2, CH4, NH3, H2CCH2, N2, Cl2, and many large biological molecules are covalent.
Organic and biological molecules consist primarily of the non-metals: carbon, oxygen, hydrogen, and nitrogen. These elements share
valence electrons to form bonds so that the outer electron energy levels of each atom are filled and have electron configurations like those
of the nearest noble gas element. Noble gases, or inert gases, are in the last column on the right of the periodic table and have complete
outer shells filled. Non-metals forming bonds with other non-metal elements tend to form covalent bonds. Non-metals bonded together
form covalently bonded molecules.
For example:
Nitrogen has one lone pair and three unpaired electrons and therefore can form covalent bonds with three hydrogen
atoms to make four electron pairs around the nitrogen.
Carbon has four unpaired electrons and combines with hydrogen, nitrogen, and oxygen to
form covalent bonds sharing electron pairs.
Diatomics: H2, N2, F2, O2, I2, Cl2, Br2 (Mnemonic for memorization: Have No Fear Of Ice Cold Beverages or the
diatomic 7)
1.
What is a covalent bond?
2.
What kinds of elements does it form between?
3.
Would Strontium (Sr) form a covalent bond with Chlorine (Cl)? Why or why not?
4.
Would Carbon (C) form a covalent bond with Neon (Ne)? Explain why or why not.
5.
Would Carbon (C) form a covalent bond with Oxygen (O)? Why or Why not?
6.
Show how Chlorine would bond to Chlorine what type of bond would be formed.
7.
Show how Carbon would bond to Hydrogen to form methane (CH4).
How to draw Lewis structures for molecules that contain no charged atoms
1) Count the total valence electrons for the molecule: To do this, find the number of valence electrons for each atom in the molecule,
and add them up.
2) Figure out how many octet electrons the molecule should have, using the octet rule: The octet rule tells us that all atoms want eight
valence electrons (except for hydrogen, which wants only two), so they can be like the nearest noble gas. Use the octet rule to figure
out how many electrons each atom in the molecule should have, and add them up. The only weird element is boron - it wants six
electrons.
3) Subtract the valence electrons from octet electrons: Or, in other words, subtract the number you found in #1 above from the
number you found in #2 above. The answer you get will be equal to the number of bonding electrons in the molecule.
4) Divide the number of bonding electrons 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:
o Hydrogen and the halogens bond once.
o The family nitrogen is in bonds three times. So does
o The family oxygen is in bonds twice.
boron.
o The family carbon is in bonds 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.
6) Find the number of lone pair (nonbonding) electrons by subtracting the bonding electrons (#3 above) from the valence electrons
(#1 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.
Other methods can be used: 1) take total number of valence electrons and divide by 2 to find how many unshared and shared
pairs, and then arrange.
Practice Naming Covalent Compounds
Prefix
Number
1. Name the following molecules.
1
SF6
2
3
C4H10
4
5
CCl4
6
7
BH3
8
9
H2
10
2. Write the following formulas from the names.
Carbon tetrabromide
Heptanitrogen nonasilicide
Antimony pentasulfide
Disulfur decafluoride
In your own words, write the steps you would take to draw dot structures for molecules. Then draw the following:
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H2S
CCl2Br2
O3
N2
H2O2
O2
Covalent Compounds Worksheet
1)
2)
3)
4)
Based on the properties of the following materials, determine whether they are made of primarily ionic
compounds or covalent compounds:
a)
telephone receiver: ______________________________________
b)
concrete: ______________________________________________
c)
gasoline: ______________________________________________
d)
candy corn: ____________________________________________
Name the following covalent compounds:
a)
SiF4 __________________________________________________
b)
N2S3 _________________________________________________
c)
HBr __________________________________________________
d)
Br2 __________________________________________________
Write the formulas for the following covalent compounds:
a)
diboron hexahydride ____________________________________
b)
nitrogen tribromide ______________________________________
c)
sulfur hexachloride ______________________________________
d)
diphosphorus pentoxide __________________________________
Write the empirical formulas for the following compounds: To find the empirical formula divide by the smallest
subscript number. This becomes the mole ratio of the elements and is represented by subscripts in the empirical
formula.
a)
C2H4O2 _______________________________________________
b)
boron trichloride ________________________________________
c)
methane ______________________________________________
d)
C6H12O6 ______________________________________________
5)
List three differences between ionic and covalent compounds:
6)
Explain why ionic compounds are formed when a metal bonds with a nonmetal but covalent compounds are
formed when two nonmetals bond.
Positive
ion
Negative
ion
Formula
Formula Mass
Molar Mass
1. Sodium Iodide
Na+
I-
NaI
22.99+126.9=149.89 g
2. Silver( I ) sulfide
Ag+
S2-
Ag2S
Name of Ionic Compound
3. Barium sulfate
4. Lithium sulfide
5. Sodium hydroxide
6. Zinc sulfate
Zn2+
7. Iron(III) phosphate
Fe3+
8. Nickel (II) hydroxide
Ni2+
9.
Cr3+
Chromium (III) oxide
10. Iron (III) sulfate
11. Copper (II) nitrate
12. Copper (II) carbonate
13. Magnesium phosphide
14. Aluminum nitrate
Compare and Contrast Covalent and Ionic Bonds
1. Define the following:
a. Electronegativity
b. Ionization Energy
c. Electron
d. Covalent Bond
e. Ionic Bond
f. Valence Shell
g. Lone Pair
h. Bonding Pair
i. Nonmetal
j. Metal
2. Using 9 of the 10 definitions you just wrote, compare and contrast covalent and ionic bonds.
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Worksheet: Introduction to Bonding1
For each of the following statements, write ionic, covalent or metallic bonds.
______________ electrons are shared
______________ electrons delocalized
______________ electrons are transferred
______________ crystal lattice
______________ luster
______________ nonconductors in the solid, molten, and dissolved state
______________ malleable and ductile
______________ high melting and boiling points
______________ volatile liquids and gases
______________ weaker forces between atoms
______________ hard—difficult to crush
What is the difference between a cation and an anion?
In a polar covalent bond, the electrons are shared ____________. In a nonpolar covalent bond,
the electrons are shared ____________.
Predict what type of bond will form between the following pairs of atoms by stating element
type (metal, metalloid, nonmetal)
F and F
Na and O
Cl and Cl
Ti and O
Fe and Fe
Mg and O
O and H
Ca and S
Ag and Ag
5.5
Modified 10/5/11 by MElizabeth
Be and Cl
1
CHEMISTRY: A Study of Matter, © 2004, GPB,
Note Taking Guide: Episode 5012
A chemical bond - forms when 2 or more atoms rearrange ____________ _____________ to
increase ________________.
ionic bond - forms when valence _____________ are _____________ from one atom to
another forming salts. Ionic bonds usually form between metal and nonmetal elements.
cation – atom __________ electrons to become ______ charged
anion - atom __________ electrons to become ______ charged
In ionic compounds the ions are arranged in a ___________ ___________.
__________ forces hold the ions together (these are electrostatic forces).
Properties of ionic compounds:
 high ____________ and ___________ points
 ___________ - not easily ______________
 _____________ electricity when ___________ or______________ because the ions
are free to __________. Does not conduct electricity when a solid
Covalent bond - ______________ are _____________, forming _______________ held
together by covalent bonds. Occurs usually with non-metal elements.
Covalent compounds have ____________ forces holding the _____________ together.
Properties of covalent compounds:
• lower ____________ and ___________ points
• Many covalent compounds are _____________ liquids or gases.
• ____________– easier to ___________
• are not ________________ of electricity
Compound
Solid Conductivity
Dissolved in Water
Conductivity
Bond Type
Table Salt (NaCl)
Para dichloro benzene
Copper Sulfate (CuSO4)
Ethanol (C2H2OH)
Sugar (sucrose)
Selzer (NaHCO3)
Electronegativity - property that tells how strong an atom’s ____________ is for an
_____________.
Since oxygen has a ______________ electronegativity than hydrogen, oxygen holds onto
shared electrons ____________, giving the oxygen a __________ negative charge and the
hydrogen a partial _____________ charge.
2
CHEMISTRY: A Study of Matter© 2004, GPB, 5.1
polar covalent bonds:
• electrons are shared _________________, creating partially charged ends or _______.
nonpolar covalent bonds:
• electrons are shared ________________ because atoms have the same electronegativities
Electonegativity difference:
Greater than or equal to 1.7
Between 1.7 and 0.3
Less than or equal to 0.3
Examples:
Type of Bond
Mg and F?
S and O?
Program 501, problem set 1: What type of bond will form between . . .
Li and Cl?
C and O?
Na and Cl?
Cl and Br?
S and H?
Al and Al
metallic bond – occurs between metal elements; electrons are ________________ (creates a
“____ of ___________”)
properties of metals:
1. _______________________________________________________________
2._______________________________________________________________
3._______________________________________________________________
4._______________________________________________________________
Summary
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The Chemistry Quiz
CR1._____ CR2._____ 1._____ 2._____ 3._____ 4._____ 5._____
Identify the type of element whether metal, metalloid, and non-metal and whether it will form a
cation, anion, or if an ionic form is usually not present and draw its atomic Lewis dot structure.
O non-metal, forms an anion (O-2)
N
H
F
Xe
Ca
S
Mg
Identify the type of elements that form the following bond types
Ionic_______________________________________________
Covalent ____________________________________________
Metallic_____________________________________________
Identify the type of covalent or ionic bond present between the following elements and draw
their Lewis Dot Structures
1. O2
2. NH3
3. F2
4. N2
5. H2O
6. Cl2
7. CaCl2
2d. Students know the atoms and molecules in liquids move in a random pattern relative to one another
because the intermolecular forces are too weak to hold the atoms or molecules in a solid form
The term intramolecular force refers to bonding between atoms in compounds: ionic, covalent, and metallic bonds
The term intermolecular forces refers to attractions between compounds such as dipole-dipole (water), hydrogen
bonds which are easily broken and formed.
In any substance at any temperature,
the forces holding the material
together are opposed by the internal
energy of particle motion, which
tends to break the substance apart. In
Notice how both the particles in a
solid and liquid are close together.
If the liquid and solid states are so
similar to each other on the molecular
scale because they experience
intermolecular forces, in what way do
they differ?
The main point of difference is the
relative movement of molecules in relation to neighboring molecules. In a solid, each molecule can vibrate and
rotate in place, but cannot change places or move past other molecules. That is, they can wiggle around, but they
can’t go anywhere. In a liquid, by contrast, molecules stay close to each other but they readily move past each other.
A human analogy is helpful here3. Imagine that each student on your campus is a molecule. When you are sitting in
a classroom, you and the other students are collectively acting as a solid. You can move about in your seat, turn
around and the like. But, you do not readily change seats with other students during the class period. The students
are solid.
Now imagine a crowded, really crowded, dance floor. You are just as close to your neighbors as you were in lecture,
if not closer, but now you are free to move about and change who you are near. The gaseous state is represented by
the students moving about between classes. Each goes his or her own way, briefly interacting but mainly moving
independently of one another.
When enough energy is added to the solid, the kinetic energy of the atoms and molecules increases sufficiently to
overcome the attractive forces between the particles, and they break free of their fixed lattice positions. This
change, called melting, forms a liquid, which is disordered and non-rigid. The particles in the liquid are free to move
about randomly although they remain in contact with each other. The particles in a gas are free to move about
randomly
INTERMOLECULAR
FORCES
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3
http://employees.oneonta.edu/viningwj/Chem112/Chapter%2011_Forces_Between_Molecules_and_the_Liquid_State_1_9.pdf
Naming Mixed Ionic and Covalent Compounds
Name the following compounds. Remember, they may be either ionic or covalent compounds, so
make sure you use the right naming method!
1)
NaF __________________________________________
2)
NF3 __________________________________________
3)
Li2O __________________________________________
4)
Al2S3 __________________________________________
5)
MgSO4 __________________________________________
6)
SiH4 __________________________________________
7)
KNO3 __________________________________________
8)
P2O5 __________________________________________
9)
CH4 __________________________________________
10)
Ca(OH)2 __________________________________________
Write the formulas for the following compounds. Remember, they may be either ionic or covalent
compounds, so make sure you use the right method!
11)
lithium chloride __________________________________________
12)
nitrogen trichloride __________________________________________
13)
sodium oxide __________________________________________
14)
dinitrogen trioxide __________________________________________
15)
ammonia __________________________________________
16)
diboron dihydride __________________________________________
17)
potassium phosphide _________________________________________
18)
oxygen difluoride __________________________________________
19)
magnesium nitrate __________________________________________
20)
aluminum carbonate __________________________________________
What type of elements give
off valence electrons to
achieve a configuration of a
noble gas?
How is a covalent
bond formed?
Write the Lewis structure
for NH3
What type of elements receive
valence electrons to achieve a
configuration of a noble gas?
How is an ionic
bond formed?
Write the Lewis structure
for CO2
Which group of metals is the
most chemically active?
Write the Lewis structure for
F2
Write the Lewis structure
for C2H4
Which group of nonmetals is
the most chemically active?
Write the Lewis structure for
O2
The composition of a compound contains only C, O, and
H atoms. What type of bond
does the compound have?
According to the periodic
table, which two elements in
each set of elements will likely
form an ionic bond?
Write the Lewis structure for
H2O
Given the chemical formula of
a compound, how will you tell
whether the compound is an
ionic compound or molecular
compound?
Of the seven diatomic
molecules, which ones contains
A double bond?
Write the Lewis structure
for N2
A
1
14
17
18
B
2
3
17
18
Elements in the same group
have similar physical and
chemical properties because
they have the same number of
_____________?
What type of elements have
atoms arranged in crystal
structure?
Solid, liquid or gas?
A triple bond?
Which groups on the periodic
table would form ions with
the following charges:
a) 1+
c. 2+
b) 1-
d. 2-
Show how you determine the
mass of one mole of CO2
In the periodic table, what
group of elements has the
most stable valence electron
configuration in the ground
state?
Using the electronegativity
Show how you determine the
values on p 263, how would
mass of one mole of
you know if the bond between
Mg (C2H3O2)2
two atoms is ionic?
In a binary ionic compound,
one of the elements is Cl.
Which of the following would
the other element most likely
to be?
Given different pairs of
atoms, how would you decide
which has the most polar
bond?
Show how you determine the
mass of one mole of
Fe(NO3)3
The Lewis structure of an
atom has eight dots. Name
the elements with the same
Lewis structure.
Given several bars of
different metals each
containing one mole.
Identify the metals with the
following approximate
masses.
Cl
I
Mg
Br
What are the seven diatomic
molecules?
a) 27 g
b) 65 g
c) 59 g
d) 40 g
Which part of the periodic
table are these located?
Which one has a larger ion
than its neutral atom?
metals __________
a. Ca
b. F
nonmetals _________
c. Zn
d. Cs
metalloids _________
What type(s) of elements
would make
a. ionic bond
b. covalent bond
CHEMICAL BONDING AND ATOMIC STRUCTURE VIDEO NOTES
Learning Objectives







Visualize chemical bonds as electrostatic forces between atoms
Compare behavior of valence electrons in metallic and covalent bondings
Recognize that the making and breaking of bonds involves energy changes
Compare the structure of atomic, ionic, and molecular crystals
Use types of bonding to account for physical and chemical properties
Appreciate the importance of the bonding capacity of carbon
Use electronegativity to account for trends in bonding in the periodic table
Before Viewing
1. What are the laws of electrostatics as applied to charged bodies.
2. What are valence electrons?
After Viewing
1. Explain why a football field was used to describe the relative size of an atom?
2. What are the three kinds of bonds?
3. What common properties of Fe and Cu can be explained by the metallic bond?
4. How can you account for differences in the conductivity of iron and copper?
5. What kind of bond would you expect between atoms of two different metals? Explain
6. What kind of bond would you expect between two different non-metals? Explain
7. Explain how the bonds transmit electricity and heat differently.
8. What does molten mean?
9. What does sublimation mean (it happens to iodine crystals)
10. In what ways can atoms of elements achieve the configuration of noble gases?
11. Use atomic structures to account for the increase in electronegativity across a period; its decrease from
the top to bottom of a group.
The Alkaline Earth Elements
Background Information
The arrangement of the elements in the periodic table is one of the most important achievements
in modern chemistry. The physical and chemical properties of elements change in a regular
pattern as you go both across the rows and down the columns of the periodic table. As a result,
when elements close to each other in a row or column are compared, they have many of the same
properties. However, when elements farther away from each other in a row or column are
compared, they have more dissimilar properties.
The elements in Group 2 of the periodic table are known as the alkaline earth elements. Like all
members of a group, or family, of elements, they have certain properties that change in a regular
pattern within the group. One of these properties is the ability to form a precipitate, or solid
substance, as a result of a chemical reaction. The precipitate cannot dissolve in water and
eventually settles to the bottom of the container. In this investigation, you will compare the abilities
of the alkaline earth elements to form precipitates as a result of a chemical reaction.
Question: Will the ability of the alkaline earth elements to form precipitates increase or decrease as
you move from top to bottom?
Hypothesis:
Materials (per group)
Safety goggles
Spot plate
Sheet of notebook paper
Procedure
1. Place the spot plate in the center of a sheet
of notebook paper so that there are 4 spots
running down and 3 spots running across.
See Figure 1.
2. Along the side of the notebook paper next to
each of the four spots, write the names of
the four alkaline earth elements that are
present in each nitrate compound listed in
the materials you are using. Write them in
the same order in which they are listed. See
Figure 1.
3. Along the top of the notebook paper next to
each of the three spots, write the names of
the three substances that are combined with
potassium in the materials you are using.
See Figure 1.







Dropper bottles of:
magnesium nitrate
calcium nitrate
strontium nitrate
barium nitrate
potassium carbonate
potassium sulfate
potassium chromate
4. Put on your safety goggles. Place 3 drops of potassium carbonate in each of the four spots
under the word "carbonate." Place 3 drops of potassium sulfate in each of the four spots
under the word "sulfate." Place potassium chromate in each of the four spots under the
word "chromate." Be very careful not to mix the liquid from one spot with the liquid from
another.
5. "Shake the dropper bottle of magnesium nitrate and place 3 drops in each of the three
spots in the row labeled "magnesium." Observe each spot carefully and record the result in
the Data Table. Repeat this procedure using the dropper bottles containing calcium nitrate,
strontium nitrate, and barium nitrate. Be very careful not to mix the liquid from one spot
with the liquid from another.
6. After recording your results, wash your spot plate thoroughly with soapy water and a brush.
Observations
DATA TABLE (2 points) (Use ppt for a precipitate and NR for no reaction.)
Alkaline earth
Metal
Carbonate
Sulfate
Chromate
Magnesium
Calcium
Strontium
Barium
Conclusion: (8 points) (In the form of a paragraph.)
1. Restate your hypothesis.
2. Was your hypothesis supported or denied? Explain.
3. Was there evidence of a chemical reaction occurring in any of the spots? Explain your
answer.
4. Which alkaline earth element formed the smallest number of precipitates?
5. Which alkaline earth element formed the greatest number of precipitates?
6. What is the relationship between the number of precipitates formed and the location of the
alkaline earth element on the periodic table?
7. If the ability of an alkaline earth element to form a precipitate is an indication of its ability to
chemically react with other substances, which is the most reactive element?
The least reactive?
8. List the alkaline earth metals in order of their chemical reactivity, starting with the most
reactive.
9. How does the order of the elements you listed in question 8 compare to their order in the
periodic table?
Application and Critical Thinking (4 points) (Answer question below.)
1. Group 1 in the periodic table is known as the alkali metals. Based on your investigation of'
the Group 2 elements, predict the comparative reactivity of the elements in Group 1 of the
periodic table.
2. If you had a solution containing a mixture of magnesium, strontium, and barium, how could
you separate the three elements (Hint: Review the information in the Data Table.)
Vocabulary and Basic Concepts and Properties
Fill in the following blanks using the work bank. Metallic Neutral Nucleus Protons
Affinity Charge Conductivity Covalent Crystal lattice Force Ionic Ionization
Malleability
substances
Lowest
1. A chemical bond in an attractive _______________________ that holds atoms together.
2. Chemical bonding is the process of atoms combining to form new __________________________.
3. Matter tends to exist in its ______________________________ energy state.
4. A(n) __________________________ bond is a bond in which one atom donates electrons to another atom.
5. When the number of protons equals the number of electrons an atom has a _________________________
charge.
6. Ions are atoms with a positive or negative _______________________________.
7. _______________________________ is the process of removing electrons from atoms to form ions.
8. Electron_________________________________ is the tendency of an atom to gain electrons when forming
bonds. Challenge: What is the term for an atom’s tendency to take electrons:_________________________
9. A bond in which atoms share electrons is called a _________________________ bond.
10. In a(n) ____________________________ bond many electrons are share by many atoms.
11. Metallic bonds are ____________________________________ thus metals are able to be pounded into many
shapes.
12. Ionic compounds have a low _____________________________ in the solid state, and a higher
_________________________(same work) in the molten state.
Indicate whether the following statements are true (T) or false (F). If the statement is false, re-write the
statement to make it true.
1. Chemical bonding is the process of atoms combining to form new substances.
2. Valence electrons are in the innermost energy level.
3. Matter in its lowest energy state tends to be more stable.
4. Particles with a positive or negative charge are called ions.
5. One property common to metals is ductibility.
6. Covalent molecules tend to have higher melting and boiling points compared to ionic compounds.
7. Covalent molecules conduct electricity in all states.
8. Hydrogen bonding intermolecular forces are stronger than London Dispersion intermolecular forces.
9. Ionic compounds typically exist in the gaseous phase at room temperature.
10. When an atom loses on or more electrons it becomes negatively charged and we call it a cation.
11. Polar molecules have a permanent dipole moment.
1. Chemical compounds are formed when atoms are bonded together



Breaking a chemical bond is an endothermic process.
Forming a chemical bond is an exothermic process.
Compounds have less potential energy than the individual atoms they are formed from.
2. Two major categories of compounds are ionic and molecular (covalent) compounds.
3. Compounds can be differentiated by their chemical and physical properties.
 Ionic substances have high melting and boiling points, form crystals, dissolve in water (dissociation), and conduct electricity
in solution and as a liquid.
 Covalent or molecular substances have lower melting and boiling points, do not conduct electricity.
 Polar substances are dissolved only by another polar substance. Non-polar substances are dissolved only by other nonpolar substances.
4. Chemical bonds are formed when valence electrons are:
 Transferred from one atom to another – ionic.
 Shared between atoms – covalent.
 Mobile in a free moving “sea” of electrons – metallic.
5. In multiple (double or triple) covalent bonds more than 1 pair of electrons are shared between two atoms.
6. Polarity of a molecule can be determined by its shape and the distribution of the charge.
 Polar molecules must have polar covalent bonds.
 Polar molecules are asymmetrical.
 Nonpolar molecules are symmetrical and/or have no polar covalent bonds.
7. When an atom gains an electron, it becomes a negative ion, anion, and its radius increases.
8. When an atom loses an electron, it becomes a positive ion, cation, and its radius decreases.
9. Atoms gain a stable electron configuration by bonding with other atoms.
 Atoms are stable when they have a full valence level.
 Most atoms need 8 electrons to fill their valence level.
 H and He only need 2 electrons to fill their valence level.
 The noble gasses (group 18) have filled valence levels. They do not normally bond with other atoms.
10. Electron-dot diagrams (Lewis structures) represent the valence electrons in elements, compounds and ions.
 Electrons in Lewis structures are arranged by their orbitals.
 The first two electrons are placed together in the “s” orbital.
 The remaining electrons are spread among the 3 “p” orbitals.
 The “s” orbital must be filled first. Then each “p” orbital must have one electron before another “p” orbital gains a second.
 The filling of electrons in a dot diagram is accomplished by putting one dot on each of four sides before doubling up.
11. Electronegativity indicates how strongly an atom of an element attracts electrons in a chemical bond. These values are
based on an arbitrary scale. Electronegativity can also be described as electroaffinity.
12. The electronegativity difference between two bonded atoms can determine the type of bond and its polarity.
0.0 - 0.4 = non-polar covalent
0.4-1.7 = polar covalent
1.7+ = ionic
13. Bonding guidelines:
 Metals react with nonmetals to form ionic compounds (salts).
 Nonmetals bond with nonmetals to form covalent compounds (molecules).
 Ionic compounds with polyatomic ions have both ionic and covalent bonds.
14. Intermolecular forces allow different particles to be attracted to each other to form solids and liquids.
 Hydrogen bonds are an example of a strong IMF between atoms.
 Hydrogen bonds exist between atoms of hydrogen and oxygen, fluorine, or nitrogen.
 Substances with hydrogen bonds tend to have much higher melting and boiling points than those without hydrogen bonds.
15. Physical properties of a substance can be explained in terms of chemical bonds and intermolecular forces. These include
conductivity, malleability, solubility, ductility, hardness, melting point and boiling point.
Unit 3 Bonding Study Guide
Chemistry Standard Set 2
Key Vocabulary Terms
1.
2.
3.
4.
5.
6.
7.
8.
Ionic Bond
Covalent Bond
Metallic Bond
Oxidation State
Coefficient
Subscript
Superscript
Metal
9. Non-metal
10. Neutrality
11. Electrostatic Attraction
12. Cation
13. Anion
14. Shared Pairs
15. Unshared Pairs
16. Outer Electrons/Valence Electron
Concepts
1.
2.
3.
4.
Orbitals – Filled – octet rule.
Energy – energy to form bonds, bond energy
Crystal Lattice – Alternating positive and negative ions to form ionic solids
Properties of Covalent, Ionic, and Metallic Compounds
Covalent – non-metals, lower melting points, non-conductor of electricity
Ionic – metal and non-metals, higher melting points, conductor of electricity
Metallic – metals, conductor of electricity and heat, malleability, ductility, reacts with acid
5. Neutrality in atoms and in the formation of formulas (sum of positive charge = sum of negative
charge)
6. Nomenclature- primarily ionic (nonmetal ending –ide) and some covalent (e.g., mono, di , and
tri)
7. Formation and Dissolving
Items for Memorization
Polyatomic Ions
Nitrate ion
Sulfate ion
Ammonium ion
Phosphate ion
Hydroxide ion
Carbonate ion
NO3-1
SO4-2
NH4+1
PO4-3
OH-1
CO3-2
Diatomic Molecules
Hydrogen
H2
Nitrogen
N2
Chlorine
Cl2
Iodine
I2
Bromine
Br2
Fluorine
F2
Oxygen
O2
Transition Metal Oxidation
States
Iron
Fe II and III
Copper
Cu I and II
Zinc
Zn II
Gold
Au I and III
Mercury
Hg I and II
Lead
Pb II and IV
Skills
1. Ability to draw Lewis Dot Structures
Atoms (e.g., Na, Ba, Al, C, N, O, Cl, and Ne)
Covalent Compound (e.g., N2, O2, Cl2, and H2)
Polyatomic Structure (e.g., H2O, CO2, and CH4)
2. Ability to inventory compounds: Number and Type of Atoms (e.g., Al2(CO3)3, ( NH4)3PO4, and
Ba(OH)2)
3. Ability to name ionic compounds given the formula and write the formula given a name (e.g.,
magnesium bromide, calcium hydroxide, iron III sulfate, and gold I phosphate)
4. Ability to predict bond types (i.e., ionic, covalent, and metallic)
Atomic Structure Review Worksheet
1. The 3 particles and respective charges of the atom are:
a. ______________________
b. ______________________
c. ______________________
2. The number of protons in one atom of an element determines the atom’s_______________________ , and
the number of electrons determines ______________________ of and element.
3. The atomic number tells you the number of ______________________ in one atom of an element. It also
tells you the number of ______________________ in a neutral atom of that element. The atomic number
gives the “identity “ of an element as well as its location on the Periodic Table. No two different
elements will have the ______________________ atomic number.
4. The _________________________________ of an element is the average mass of an element’s naturally
occurring atom, or isotopes, taking into account the ______________________ of each isotope.
5. The ____________________________________ of an element is the total number of protons and neutrons in the
______________________ of the atom.
6. The mass number is used to calculate the number of ______________________ in one atom of an element.
In order to calculate the number of neutrons you must subtract the ______________________ from the
______________________ .
7. Give the symbol and number of protons in one atom of:
Lithium __________________
Bromine __________________Iron ______________________ Copper __________________
Oxygen __________________
Mercury __________________Krypton __________________ Helium __________________
8. Give the symbol and number of electrons in a neutral atom of:
Uranium __________________ Chlorine __________________Boron __________________ Iodine __________________
Antimony __________________ Xenin __________________
9. Give the symbol and number of neutrons of the most common isotope in one atom of:
(To get “mass number”, you must round the “atomic mass” to the nearest whole number)
Show your calculations.
Barium __________________
Bismuth __________________
Carbon __________________
Hydrogen __________________
Fluorine __________________
Magnesium __________________
10. Name the element which has the following numbers of particles:
a. 26 electrons, 29 neutrons, 26 protons _____________________
b. 53 protons, 74 neutrons _____________________
c. 2 electrons (neutral atoms) _____________________
d. 20 protons _____________________
e. 86 electrons, 125 neutrons, 82 protons (charged atom) _____________________
f. 0 neutrons _____________________
11. If you know only the following information can you always determine what the element is? (Yes/No).
a. number of protons ___________
b. number of neutrons___________
c. number of electrons in a neutral atom___________
d. number of electrons___________
PRACTICE TEST
1. During a flame test, ions of a specific metal are heated in the flame of a gas burner. A characteristic color of light is
emitted by these ions in the flame when the electrons
a) gain energy as they return to lower energy levels
b) gain energy as they move to higher energy levels
c) emit energy as they return to lower energy levels
d) emit energy as they move to higher energy Levels
2. The high electrical conductivity of metals is primarily due to
a) high ionization energies
b) filled energy levels
c) mobile electrons
d) high electronegativity’s
3. Which statement best describes the density of an atom’s nucleus?
a) The nucleus occupies most of the atom’s volume but contains little of its mass.
b) The nucleus occupies very little of the atom’s volume and contains little of its mass.
c) The nucleus occupies most of the atom’s volume and contains most of its mass.
d) The nucleus occupies very little of the atom’s volume but contains most of its mass.
4. Which Group of the Periodic Table contains atoms with a stable outer electron configuration?
a) 1
b) 16
c) 8
d) 18
5. An atom of carbon-12 and an atom of carbon-14 differ in
a) atomic number
b) mass number
c) nuclear charge
d) number of electrons
6. Which list of elements contains two metalloids?
a) Si, Ge, Po, Pb
b) Si, P, S, Cl
c) As, Bi, Br, Kr
d) Po, Al, I, Xe
7. Which is the negatively charged particle located outside the nucleus of an atom?
(1)
electron
(2)
neutron
(3)
silicon
(4)
Which of the atoms pictured is not likely to form an ion?
(1) C
(2) Na
(3) O
(4) they are all equally likely to form an ion
proton
8. Which of the atoms pictured is likely to form an anion?
(1) C
(2) Na
(3) O
(4) they are all equally likely to form an ion
9. In the periodic table of the elements, elements in Group 1 have how many outer electrons?
(1) 0
(2) 1
(3) 4
(4) 8
10. Which type of nuclear radiation (beta particles, gamma rays, or alpha particles) can be blocked by…?
a) a piece of paper
b) a block of wood
c) a piece of lead
_____________
___________
____________
11. Elements from which two groups in the periodic table would most likely combine with each other to form an ionic
compound?
A. 1 and 2
B.
1 and 17
C. 16 and 17
D. 17 and 18
12. The diagram below shows the structure of a brain chemical called acetylcholine:
Based on the nature of the elements making up acetylcholine, the bonds present in the compound are most likely…
A. nuclear
C. hydrogen
B. metallic
D. covalent
13. Which of the following elements can form an anion that contains 54 electrons, 74 neutrons, and 53 protons?
The illustration below shows two atoms of a fictitious element (M) forming a diatomic molecule.
14. What type of bonding occurs between these two atoms?
A. covalent
B. ionic
C. nuclear
D. polar
15. When elements from group 1 (1A) combine with elements from group 17 (7A), they produce compounds. Which of
the following the correct combining ratio is between group 1 (1A) elements and group 17 (7A) elements?
A. 1:1
B. 1:2
C. 2:1
D. 3:2
16. The bonds in BaO are best described asA. covalent, because valence electrons are shared
B. covalent, because valence electrons are transferred
C. ionic, because valence electrons are shared
D. ionic, because valence electrons are transferred
17. The strength of an atom’s attraction for the electrons in a chemical bond is the atom’sA. electronegativity
C. heat of reaction
B. ionization energy
D. heat of formation
18. The chemical bond between which two atoms is most polar?
A. C–N
B. H–H
C. S–Cl
D. Si–O
19. When cations and anions join, they form what kind of chemical bond?
A. ionic
B. hydrogen C. metallic
D. covalent
20. Which of the following atoms has six valence electrons?
A. magnesium (Mg)
B. silicon (Si)
C. sulfur (S)
D. argon (Ar)
21. Which change in state would involve a decrease in the intermolecular force of attraction holding the water
particles together?
A. H2O(l) → H2O(s)
B. H2O(g) → H2O(l)
C. H2O(g) → H2O(s)
D. H2O(s) → H2O(l)