Lesson 1: Intro to Bonding & Types of Bonds

PRACTICE PACKET LEVEL 5: BONDING

Regents Chemistry: Mr. Palermo

Practice Packet

Level 5: Bonding

LESSON 1: INTRO TO BONDING & TYPES OF BONDS

For each phrase, check either “bond breaking” or “bond forming”

Bond Breaking / Bond Forming
Stability of the chemical system increases
Energy is released
b. / Cl + Cl  Cl2
c. / exothermic
d. / endothermic
e. / N2 N + N
f. / Energy is absorbed
g. / Stability of the chemical system decreases
Substance / Bond type
a. / NaCl(s)
b. / CO2 (g)
c. / NO (g)
d. / Cu (s)
e. / MgBr2 (g)
g. / HCl (aq)
h. / SO2 (g)
i. / AlCl3 (s)
j. / Ag (s)
k. / NaBr(s)

2 Identify which bond type is described by each statement below. Choices: ionic, covalent, metallic

For each example provide the molecule, bond and determine when and if it conducts electricity

Type of Molecule
(metallic, ionic, molecular) / Type of Bond
(Metallic, ionic, covalent, both ionic and covalent) / Conducts electricity?
(check all that apply)
No (s) (l) (aq)
a.Li2O
b.AlCl3
c.F2
d.CH4
e.HI
f.Fe
g.Na3PO4
h.CaO
i.C(diamond)
j.C(graphite)
k.H2
l.Na
m.NH4Br
n.KNO3
o.O3
p.SiO2
q.NH3
r.FeBr2
s.Hg
t.CO2

Indicate which type of substance is described by each statement.

Choices: covalent (molecular), ionic, metallic

Type of substance
a. / Can conduct electricity in the solid and liquid phases
b. / A soft substance whose atoms are held together by
covalent bonds
c. / Low melting point and poor electrical conductor
d. / Can conduct electricity when aqueous or molten (liquid)

Lesson 2: Bond Polarity

Electronegativity values generally ______down a group and ______across a period.
Metals tend to have ______electronegativity values and nonmetals are ______values.

Using your table above find the electronegativity difference for each substance.

Then, check which bonds are present. If it’s a metal and a nonmetal it is automatically Ionic.

Substance / Electronegativity difference(s) / Ionic / Covalent / Polar / Nonpolar
I2
PCl3
SiO2
Br2
CO2
NaCl
CH4
N2O5
NH3
KCl

Indicate which atom will have the positive charge and which will have the negative charge in the following polar bonds:

H-ClH-FS-FN-O

What is electronegativity?

What factor causes some combinations or atoms to form ions, and other combinations of atoms to form covalent bonds. Explain in detail.

What is a nonpolar covalent bond? Explain the electronegativity differences attributed to this type of bond.

What is a polar bond? Explain the electronegativity differences attributed to this type of bond.

Explain the relationship between electronegativity and polarity.

What is a dipole?

What symbol indicates a partial charge? ______

How do you determine which atom gets the partial negative charge?

Given the following indicate which atom will receive the partial negative charge and which atom will receive the partial positive charge. Place the partial charges in the upper right hand corner of the atom symbol:
H – Cl
H – F
S –F
N – O

Compare the degree (which compound is most polar, which is least polar) of polarity in HF, HBr, HCl, and HI.

Classify the type of molecule the diagrams below represent (Ionic, Polar Covalent, or Nonpolar Covalent), and explain your reasoning.

Electron Distribution Diagram / Type of Compound / Reason for Classification of Compound

LESSON 3: LEWIS (ELECTRON) DOT DIAGRAMS FOR IONIC COMPOUNDS

Ion / Electron-dot Diagram / Electron Configuration / Ion / Electron-dot structure / Electron Configuration
a. / sodium
Na+ / e. / oxide
O2
b. / aluminum
Al3+ / f. / bromide
Br
c. / calcium
Ca2+ / g. / phosphide
P3
d. / magnesium
Mg2+ / h. / sulfide
S2

Complete the table below (electron dot diagrams for ions)

Ionic compound
(name& formula) / Electron-dot Diagram / Total # of ions / Ionic compound
(name& formula) / Electron-dot structure / Total # of ions
a. / sodium fluoride
NaF / f. / aluminum chloride
b. / potassium chloride
KCl / g. / sodium sulfide
c. / calcium iodide
CaI2 / h. / lithium hydride
d. / magnesium oxide
MgO / i. / aluminum oxide
e. / rubidium oxide
Ru2O / j. / calcium phosphide

Complete the table below (electron dot diagrams for ionic compounds)

LESSON 4: WRITING FORMULAS FOR IONIC COMPOUNDS

Name / Criss-Cross, Reduce / Formula / Electron-dot diagram
potassium fluoride
lithium bromide
strontium chloride
barium iodide
gallium nitride
zinc sulfide

Complete the table for the following ionic substances

Write the formula for the ionic substances below containing polyatomics (use table E)

Name / Criss-Cross, Reduce / Formula / Name / Criss-Cross, Reduce / Formula
sodium sulfate / barium phosphate
aluminum chromate / calcium hydroxide
magnesium hydrogen carbonate / potassium hydrogen sulfate
lithium permanganate / ammoniumchloride
rubidium oxalate / sodium acetate

More Practice Writing Ionic Formulas

Write the formula for each ionic compound below. Remember to Criss cross and reduce the oxidation states. The first problem is done for you.

Name / Cation (+) / Anion (-) / Criss Cross & Reduce / Formula
1 / Sodium Chloride / Na1+ / Cl1- /
2 / Aluminum Chloride / Al3+ / Cl1-
3 / Aluminum Phosphide
4 / Magnesium Oxide
5 / Cesium Fluoride
6 / Strontium Nitride
7 / Lithium Sulfide
8 / Calcium Chloride
9 / Sodium Bromide
10 / Beryllium Iodide
11 / Strontium Fluoride
12 / Aluminum Fluoride
13 / Potassium Nitride
14 / Sodium Sulfide
15 / Lithium Oxide

2. Writing Formulas containing Polyatomic Ions

Ionic Compound

/ Criss Cross and Reduce / Formula

1. calcium carbonate

barium nitrate

ammonium sulfate

aluminum hydroxide

calcium phosphate

cesium nitrate

sodium nitrite

8. calcium sulfate

beryllium sulfate

sodium carbonate

magnesium phosphate

14. ammonium nitrate

3. What is the sum of the charges on all ionic compounds?

LESSON 5: NAMING IONIC COMPOUNDS

Formula / How MANY oxidation states listed for the Metal? / Name
One / Two or more*
LiBr
Ag2O
SnO / tin (II) oxide
Ba3N2
AgBr
Cu3P
Mg(NO3)2
Co2O3
NaNO3
KI
NaClO
Fe(OH)3
PbSO4
NaHCO3
Ni2(SO4)3
Ti2O3
Al2(SO3)3
Al(CN)3
NH4Cl
KNO3
CaCO3
(NH4)2CO3

MORE NAMING PRACTICE:

Formula / How MANY oxidation states listed for the Metal? / Name
One / Two or more*

AlF3

FeCl2

Co(NO3)2

MgSO4

NH4Cl

CrPO4

Ba(OH)2

Na2CO3

BaF2

Cu(NO3)2

NiSO4

Zn3(PO4)2

Na3N

LESSON 6: NAMING AND FORMULA WRITING: COVALENT (MOLECULAR) COMPOUNDS

Name / Formula / Name / Formula
dinitrogen trioxide / silicon tetrafluoride
diphosphoruspentoxide / carbon tetrachloride
sulfur dioxide / boron triiodide
silicon dioxide / carbon disulfide
xenon pentafluoride / phosphorus pentabromide
dihydrogen monoxide / boron trihydride

1.Write formulas for the following molecular substances.

Name / Formula / Name / Formula
N2O5 / H2S
SF5 / BF3
PBr3 / PH3
SO3 / H2O
B2H4 / Cl2
PCl5 / PCl3
P2O5 / SCl6
CS2 / CO2
CO / NO
BCl3 / NO2

2.Write IUPAC Names for the following molecular (covalent) substances

LESSON 7: LEWIS (ELECTRON) DOT DIAGRAMS FOR COVALENT SUBSTANCES

Complete the chart.

Molecule
(name & formula) / Total # of valence
e-’s / Electron-dot structure / Molecule
(name & formula) / Total # of valence e-‘s / Electron-dot structure
a. / methane
CH4 / g. / carbon tetrachloride
CCl4
b. / nitrogen
N2 / h. / carbon dioxide
CO2
c. / ammonia
NH3 / i. / phosphorus trichloride
PCl3
d. / water
H2O / j. / dihydrogenmonosulfide
H2S
e. / oxygen
O2 / k. / carbon monoxide
CO
f. / fluorine
F2 / l. / hydrogen
H2

MORE PRACTICE:

Draw the lewis (electron) dot structures for the following molecular (covalent) substances.

Cl2 / SH2
H2 / CF4
CS2 / SiO2
SF2 / HF

LESSON 8: MOLECULAR POLARITY

Fill in the chart below.

In terms of lone pair electrons, how can you determine if a molecule is polar?

What molecular shapes are always polar?

What molecular shapes are always nonpolar?

How can a molecule be nonpolar if it contains polar bonds?

Fill in the chart below

Molecule / Dot Diagram / Distribution of Charge (symmetrical or asymmetrical) / Molecular Polarity (polar or nonpolar molecule) / Molecular Shape (linear, pyramidal, tetrahedral or bent)
CCl4
NF3
Br2
CS2
SiO2

LESSON 9: INTERMOLECULAR FORCES (IMF’S)

Which of the following will have the higher boiling point? Explain your answer using intermolecular forces. NH3 or N2

Why does dry ice (solid CO2) evaporate before sodium chloride?

Why does gasoline (C8H18) exist in the liquid form while methane (CH4), the gas we use to power out bunsen burners, exists in the gas form even though both compounds are nonpolar?

Identify the intermolecular forces that exist in the following molecules.

Compound / Type of IMF
H2O
N2
HCl
LiCl

Of the compounds in question 4, which has the strongest surface tension?

In terms of the forces of attraction holding them together, explain why a NaCl crystal has a melting point of 800C while an ice cube of pure water has a melting point of 0C.

List the noble gases from highest to lowest boiling point. Explain your answer based on intermolecular forces of attraction.

Explain why I2 is a solid, Br2 is a liquid but Cl2and F2 are gases even though they are all Halogens.

List the following substances from highest to lowest melting point; use attractive force to justify your answers. KCl, Cl2, CH4, H2O, PCl3

11.

Network Solids

The best example of a network solid is a “diamond”. Look at the model of a diamond below. Note that the carbon atoms are bonded together with covalent bonds. The basic building unit is an atom of carbon. The structure has a very definite tetrahedral crystal shape, because these atoms are arranged and held rigidly in a fixed pattern. A diamond is very hard (a “10” on the Moh’s Scale of Hardness…the highest value possible). In order to scratch a diamond, you must break 1000’s of very strong covalent bonds! Similarly, to melt (or boil) a network solid, like a diamond, you must break 1000’s of these covalent bonds. This involves considerable energy and is the reason for their high melting points. It is because of these high temperatures and their hardness that network solids are frequently used in industry as “abrasives” (on sandpaper and on the tips of drills for cutting tools). You don’t have to worry about them melting if it gets too hot from friction or being scratched and dulled when contacting most other surfaces. Network solids have the type of properties you would expect from atoms being held together via strong covalent bonds, e.g. diamonds. They have very high melting points and are practically insoluble; are mostly nonconductors (no free electrons or ions); and they are very brittle (atoms must maintain a fixed crystal structure, if they are pushed too close together …. they repel).

Graphite is also shown below. Note that it is also pure carbon, like a diamond. However, the covalent bonds only attach carbon atoms in 2 directions, not 3 like diamonds. The dashed lines between the layers of covalently bonded carbon atoms represent weak Van der Waal forces. Graphite is a 2 dimensional network solid. The strong covalent bonds only go in “plates”, in 2 directions. The “plates” are connected via weak VDW forces. Graphite STILL has a high melting point. – You must weaken/break all of its bonds (VDW and Covalent) to melt it. But since the weaker VDW forces are present and break easily, graphite is often used as a “dry lubricant”. If you squirt graphite dust into a lock, it will lodge between the lock’s moving metallic parts. When you put in a key and turn, the graphite structure will break apart between its “plates.” VDW’s break and make it turn more easily. Graphite also has free, “delocalized” electrons (it is a resonant structure… there really are no double bonds present, but free electrons) thus…graphite is a network solid that is capable of conducting electricity. This is not characteristic of most network solids. Of course, pencil lead is graphite. What bonds break when you write??? Silicon bonds like carbon to form a network structure. The computer industry depends on “silicon chips” ,which are made conductive by placing impurities in their structure. These then provide for free electrons and allow the chip to do its job. But, pure silicon does not conduct.

Many network solids are composed of various combinations of relatively few elements on the periodic table. The elements B, C, Al, Si are found in many network solids. They can be pure or combine with one another or combine with elements near them. For example, SiO2, quartz is an example of a network solid. Corundum, Al2O3, is a network and a common abrasive used on “sandpaper”. Many gem stones, like diamond, are network solids. Emerald is made of the mineral “beryl”. Its formula is Be3Al2 (Si6O18) . Ruby is a form of corundum.

1. What is a Network solid? Give an example

2. What are some physical properties of Network solids?

Unit Review/Study Guide

INTRODUCTION TO BONDING

Elements are the simplest form of matter and cannot be decomposed. Compounds can be formed between two or more elements. They can be decomposed chemically.

a. Which of the following is a compound? NeH2OBeF

b. Which of the following cannot be decomposed by chemical means?

C12H24NH3LiCS2

Atoms bond in order to obtain a stable electron configuration, like noble gases, called the octet. Most atoms will gain or lose electrons in order to have eight valence electrons. However, small elements such as H, Li, and Be will settle for two valence electrons. Obtaining an octet makes the atoms more stable and they can release energy. The electrons obtain the octet by sharing or transferring electrons.

a. Draw the Lewis dot diagram of the following elements:

NaMgAlSi

PSClAr

b. Draw the Lewis dot diagram of the following ions:

Na+Mg+2Al+3

P-3S-2Cl-

c. Explain why the metals lost electrons but the nonmetals gained electrons.

______

d. Fill the blanks with release or absorb: “When atoms bond they ______energy. In order to break a bond, energy must be ______.

IONIC BONDING

Compounds that form between a metal and a nonmetal contain ionic bonds, transferring electrons. Ionic bonds are strong. Ionic compounds have high melting points, are generally solids at room temperature, and conduct in the liquid phase.

a. Which of the following has ionic bonds?NaClNH3Mg

b. Which of the following transfers electrons?MgBr2LiCO2

c. Which of the following has a higher melting point?CuC6H12LiF

d. Which of the following can conduct in the aqueous phase? NOKINe

COVALENT BONDING

Compounds that form between two nonmetals have covalent bonds, sharing electrons. Covalent bonds are weaker than ionic bonds. Covalent compounds have low melting points, are generally gases, liquids, or powdery solids at room temperature, and never conduct. These are also known as molecular compounds.

a. Which of the following has covalent bonds?HFLiClRb

b. Which of the following shares electrons?H2OAgCaCl2

c. Which of the following can never conduct electricity? KrRb2OH2O

d. Which of the following has both ionic and covalent bonds? LiNH3CaCO3

e. Which of the following is a molecular compound?H2OMgLiBr

METALLIC BONDING

Metallic Bonds form when a metal loses their valence electrons and a “sea of mobile electrons” form that allows the metal to conduct electricity in the solid or liquid phase.

a. Which of the following is metallic?NaClNH3Mg

b. Which of the following has a sea of mobile electrons?CuC6H12LiF

c. Which of the following can conduct in the solid phase?NeAgCaCl2

NAMING COMPOUNDS/FORMULA WRITING

When Ionic Compounds, always name the positive, cation first and then the negative, anion last. The elements are named in the same order they appear on the periodic table.When compounds have more than 2 elements, it contains a polyatomic ion. Use Table E on page 2 of your reference tables. Transition Metals are in the middle group of the periodic table. Nonmetals are on the right side of the staircase. They have multiple charges or oxidation numbers and so you must show which charge you are using with roman numerals. Polyatomic ions are a group of 2 or more atoms that are bonded very strongly and act as one ion. Name the following:

CaCl2NaFLiOHKNO3

CuBr2CuBr3Ni(OH)2NiCl3

To write a formula of an ionic compound, write the two ions separately showing their charges. Charges are on the periodic table. Then, swap the two numbers and drop the sign. Write the formula for the following:

Sodium fluorideCesium oxide

Strontium acetateAluminum phosphate

Iron(III) iodideManganese (VII) oxide

When naming Covalent Compounds, use prefixes to indicate the number of each atom present in the compound. Determine the prefix of each element using the subscript #. Remember, if only 1 atom is present for the first element do not use the prefix mono for that atom. Name the following:

HClPCl5N2O2NH3

To write a formula of a covalent compound, write the least electronegative element first. Determine the prefix of each element using the subscript #. Write the formula for the following:

Sulfur hexafluorideNitrogen dioxide

Carbon DioxideNitrogen monoxide

LEWIS STRUCTURES/GEOMETRY

Ionic Lewis diagrams show the ions involve in the bond, but no arrangement. Covalent Lewis diagrams show the sharing of electrons with lines representing two electrons. They form shapes such as linear, bent, pyramidal, and tetrahedral.

Draw the following and give the number of shared pairs, unshared pairs, and the shape if applicable.

LiFNH3

MgF2CH4

Cl2H2O

POLARITY

Bonds are polar when two atoms have different electronegativities and share unevenly. The more electronegative atom has the electrons more of the time. Nonpolar bonds form when two atoms have the same electronegativity values and share equally.

Label the bonds as polar or nonpolar:

NH3CH4Cl2H2O

Molecules are polar when the molecule is asymmetrical. They are nonpolar if the molecule is symmetrical.

Label the bonds as polar or nonpolar (Use your drawing to help you):

NH3CH4Cl2H2O

INTERMOLECULAR FORCES

Intermolecular forces are what keeps molecules together (not atoms-that’s bonds) and are responsible for phases, phase changes, surface tension and various other properties. Nonpolar molecules have the weakest attractive forces dependent on their size (the bigger the stronger). Polar molecules have stronger forces dependent on their polarity. Hydrogen bonds are a special case of polar forces between H and either F,O, or N. Molecules that are hydrogen bonded have high melting and boiling points, strong surface tension, and have closely packed particles.

Which of the following has the highest melting point? ______

HFHClHBrHI

Which of the above has the lowest boiling point? ______