PRACTICE PACKET LEVEL 5: BONDING
Regents Chemistry: Mr. Palermo
Practice Packet Level 5: Bonding
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LESSON 1: INTRO TO BONDING & TYPES OF BONDS
1. For each phrase, check either “bond breaking” or “bond forming” Bond Bond Breaking Forming
1. Stability of the chemical system increases
2. 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)
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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
3. For each example provide the molecule, bond and determine when and if it conducts electricity
Type of Molecule Type of Bond Conducts (metallic, ionic, (Metallic, ionic, covalent, electricity? molecular) both ionic and covalent) (check all that apply)
No (s) (l) (aq) Li O 2 b. AlCl 3 c. F 2 d. CH 4 e. HI f. Fe 3 www.mrpalermo.com PRACTICE PACKET LEVEL 5: BONDING
g. Na PO 3 4 h. CaO i. C (diamond) j. C (graphite) k. H 2 l. Na m. NH Br 4 n. KNO 3 o. O 3 p. SiO 2 q. NH 3 r. FeBr 2 s. Hg t. CO 2
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
A soft substance whose atoms are held together by b. covalent bonds c. Low melting point and poor electrical conductor d. Can conduct electricity when aqueous or molten (liquid) Lesson 2: Bond Polarity
1. Electronegativity values generally ______down a group and ______across a period. 4 www.mrpalermo.com PRACTICE PACKET LEVEL 5: BONDING
2. 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 Ionic Covalent Polar Nonpolar difference(s)
I2
PCl3
SiO2
Br2
CO2
NaCl
CH4
N2O5
NH3
KCl
3. Indicate which atom will have the positive charge and which will have the negative charge in the following polar bonds:
H-Cl H-F S-F N-O
4. What is electronegativity?
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5. What factor causes some combinations or atoms to form ions, and other combinations of atoms to form covalent bonds. Explain in detail.
6. What is a nonpolar covalent bond? Explain the electronegativity differences attributed to this type of bond.
7. What is a polar bond? Explain the electronegativity differences attributed to this type of bond.
8. Explain the relationship between electronegativity and polarity.
9. What is a dipole?
10. What symbol indicates a partial charge? ______
11. How do you determine which atom gets the partial negative charge?
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12. 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:
a. H – Cl
b. H – F
c. S –F
d. N – O
13. Compare the degree (which compound is most polar, which is least polar) of polarity in HF, HBr, HCl, and HI.
14. Classify the type of molecule the diagrams below represent (Ionic, Polar Covalent, or Nonpolar Covalent), and explain your reasoning.
Electron Distribution Type of Compound Reason for Classification of Compound Diagram
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LESSON 3: LEWIS (ELECTRON) DOT DIAGRAMS FOR IONIC COMPOUNDS
Electron- Electron Electron- Electron Ion dot Configuration Ion dot Configuration Diagram structure
sodium oxide a. e. Na+ O2
aluminum bromide b. f. Al3+ Br
calcium phosphide c. g. Ca2+ P3
magnesium sulfide d. h. Mg2+ S2
1. Complete the table below (electron dot diagrams for ions)
Ionic Electron-dot Total Ionic Electron-dot Total compound Diagram # of compound structure # of ions ions (name & formula) (name & formula)
sodium aluminum a. fluoride f. chloride NaF
potassium b. chloride g. sodium sulfide KCl c. calcium h. lithium
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iodide hydride
CaI2
magnesium aluminum d. oxide i. oxide MgO
rubidium calcium e. oxide j. phosphide
Ru2O
2. Complete the table below (electron dot diagrams for ionic compounds) LESSON 4: WRITING FORMULAS FOR IONIC COMPOUNDS
Name Criss-Cross, Formula Electron-dot diagram Reduce
potassium fluoride
lithium bromide
strontium chloride
barium iodide
gallium nitride
zinc sulfide
1. Complete the table for the following ionic substances
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2. Write the formula for the ionic substances below containing polyatomics (use table E)
Criss-Cross, Criss-Cross, Name Formula Name Formula Reduce Reduce
barium sodium sulfate phosphate aluminum calcium chromate hydroxide magnesium potassium hydrogen hydrogen carbonate sulfate lithium ammonium permanganate chloride
sodium rubidium oxalate acetate
More Practice Writing Ionic Formulas
1. Write the formula for each ionic compound below. Remember to Criss cross and reduce the oxidation states. The first problem is done for you.
Criss Cross & Formula Name Cation (+) Anion (-) Reduce
1 Sodium Chloride Na1+ Cl1-
Aluminum 2 3+ 1- Chloride Al Cl
Aluminum 3 Phosphide
4 Magnesium Oxide
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5 Cesium Fluoride
6 Strontium Nitride
7 Lithium Sulfide
8 Calcium Chloride
9 Sodium Bromide
10 Beryllium Iodide
11 Strontium Fluoride
Aluminum 12 Fluoride
13 Potassium Nitride
14 Sodium Sulfide
15 Lithium Oxide
2. Writing Formulas containing Polyatomic Ions
1. Ionic Compound Criss Cross and Reduce Formula
2. 1. calcium carbonate
3. barium nitrate
4. ammonium sulfate
5. aluminum hydroxide
6. calcium phosphate
7. cesium nitrate
8. sodium nitrite
8. calcium sulfate
10. beryllium sulfate
11. sodium carbonate
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12. magnesium
phosphate 14. ammonium nitrate
3. What is the sum of the charges on all ionic compounds?
LESSON 5: NAMING IONIC COMPOUNDS
How MANY oxidation states listed for the Metal?
Formula One Two or more* Name
LiBr
Ag2O
SnO tin (II) oxide
Ba3N2
AgBr
Cu3P
Mg(NO3)2
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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:
How MANY oxidation states listed for the Metal?
Formula One Two or more* Name
1. KOH
2. LiI
3. AlF3
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4. FeCl2
5. MgO
6. Co(NO3)2
7. MgSO4
8. NH4Cl
9. CrPO4
10. Ba(OH)2
11. PbS
12. Na2CO3
13. BaF2
14. Cu(NO3)2
15. AgI
16. NiSO4
17. Zn3(PO4)2
18. Na3N
LESSON 6: NAMING AND FORMULA WRITING: COVALENT (MOLECULAR) COMPOUNDS
Name Formula Name Formula dinitrogen trioxide silicon tetrafluoride diphosphorus carbon tetrachloride pentoxide sulfur dioxide boron triiodide silicon dioxide carbon disulfide xenon pentafluoride phosphorus
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pentabromide dihydrogen monoxide boron trihydride
1. Write formulas for the following molecular substances.
Formul Name Name Formula a
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
1. Complete the chart.
Total # Electron-dot Total # Electron-dot Molecule Molecule of structure of structure (name & valence (name & valence formula) formula) e-‘s e-’s a. methane g.
CH4 carbon
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tetrachloride
CCl4
carbon b. nitrogen h. dioxide
CO2 N2
phosphorus ammonia c. i. trichloride NH3 PCl3
dihydrogen water d. j. monosulfide H2O H2S
carbon oxygen monoxide e. k. CO O2
fluorine hydrogen f. l. F2 H2
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MORE PRACTICE:
Draw the lewis (electron) dot structures for the following molecular (covalent) substances.
Cl2 SH2
H2 CF4
CS2 SiO2
SF2 HF
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LESSON 8: MOLECULAR POLARITY
1. Fill in the chart below.
2. In terms of lone pair electrons, how can you determine if a molecule is polar?
3. What molecular shapes are always polar?
4. What molecular shapes are always nonpolar?
5. How can a molecule be nonpolar if it contains polar bonds?
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Molecule Dot Diagram Distribution of Molecular Molecular Charge Polarity (polar Shape (linear, (symmetrical or or nonpolar pyramidal, asymmetrical) molecule) tetrahedral or bent)
CCl4
NF3
Br2
CS2
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SiO2
LESSON 9: INTERMOLECULAR FORCES (IMF’S) a.1.Which of the following will have the higher boiling point? Explain your answer using
intermolecular forces. NH3 or N2
a.2.Why does dry ice (solid CO2) evaporate before sodium chloride?
a.3.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?
a.4.Identify the intermolecular forces that exist in the following molecules.
Compound Type of IMF
H2O
N2
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HCl
LiCl a.5.Of the compounds in question 4, which has the strongest surface tension?
6. In terms of the forces of attraction holding them together, explain why a NaCl crystal has a melting point of 800C while an ice cube of pure water has a melting point of 0C.
7. List the noble gases from highest to lowest boiling point. Explain your answer based on intermolecular forces of attraction.
8. Explain why I2 is a solid, Br2 is a liquid but Cl2 and F2 are gases even though they are all Halogens.
9. List the following substances from highest to lowest melting point; use attractive force
to justify your answers. KCl, Cl2, CH4, H2O, PCl3
10.
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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).
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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
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“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? Ne H2O Be F b. Which of the following cannot be decomposed by chemical means?
C12H24 NH3 Li CS2
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:
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Na Mg Al Si
P S Cl Ar b. Draw the Lewis dot diagram of the following ions:
Na+ Mg+2 Al+3
P-3 S-2 Cl- 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? NaCl NH3 Mg
b. Which of the following transfers electrons? MgBr2 Li CO2
c. Which of the following has a higher melting point? Cu C6H12 LiF d. Which of the following can conduct in the aqueous phase? NO KI Ne
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.
25 www.mrpalermo.com PRACTICE PACKET LEVEL 5: BONDING a. Which of the following has covalent bonds? HF LiCl Rb
b. Which of the following shares electrons? H2O Ag CaCl2
c. Which of the following can never conduct electricity? Kr Rb2O H2O
d. Which of the following has both ionic and covalent bonds? Li NH3 CaCO3
e. Which of the following is a molecular compound? H2O Mg LiBr
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? NaCl NH3 Mg
b. Which of the following has a sea of mobile electrons? Cu C6H12 LiF
c. Which of the following can conduct in the solid phase? Ne Ag CaCl2 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:
CaCl2 NaF LiOH KNO3
CuBr2 CuBr3 Ni(OH)2 NiCl3
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 fluoride Cesium oxide
Strontium acetate Aluminum phosphate 26 www.mrpalermo.com PRACTICE PACKET LEVEL 5: BONDING
Iron(III) iodide Manganese (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:
HCl PCl5 N2O2 NH3
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 hexafluoride Nitrogen dioxide Carbon Dioxide Nitrogen 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.
a. Draw the following and give the number of shared pairs, unshared pairs, and the shape if applicable.
LiF NH3
MgF2 CH4
Cl2 H2O
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.
a. Label the bonds as polar or nonpolar:
NH3 CH4 Cl2 H2O
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Molecules are polar when the molecule is asymmetrical. They are nonpolar if the molecule is symmetrical.
b. Label the bonds as polar or nonpolar (Use your drawing to help you):
NH3 CH4 Cl2 H2O
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.
a. Which of the following has the highest melting point? ______
HF HCl HBr HI
b. Which of the above has the lowest boiling point? ______
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