CH.4 - CHEMICAL BONDING: UNDERSTANDING CLIMATE CHANGElightcat-files.s3.amazonaws.com/packets/admin... · • Generally, the reason is that ionic bonding _____ the potential energy

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CHEMISTRY - KIRSS 2E

CH.4 - CHEMICAL BONDING: UNDERSTANDING CLIMATE CHANGE

CONCEPT: ATOMIC PROPERTIES AND CHEMICAL BONDS

Before we examine the types of chemical bonding, we should ask why atoms bond at all.

• Generally, the reason is that ionic bonding ____________ the potential energy between positive and negative ions.

• Generally, the reason covalent bonds form is to follow the ____________ rule, in which the element is then

surrounded by 8 valence electrons.

There are three models of chemical bonding:

In ____________________ bonding, metals connect to non-metals.

• __________________ transfers an electron to the ________________ , creating ions with opposite charges that

are attracted to each other.

Li F Li F Li F

In _______________ bonding, non-metals connect to non-metals.

• In it the nonmetals __________________ electron pairs between their nuclei.

ClCl

In _______________ bonding, metal atoms “pool” their valence electrons to form an electron “sea” that holds the metal-ion

together



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CONCEPT: CHEMICAL BONDS (PRACTICE)

EXAMPLE: Describe each of the following as either a(n): atomic element, molecular element, molecular compound or ionic compound.

atomic element ––

molecular element ––

molecular compound ––

ionic compound ––

a. Iodine

b. NH3

c. Graphite

d. Na3P

e. Ag2(SO4)2



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CONCEPT: THE IONIC-BONDING MODEL

The central idea of ionic bonding is that the metal transfers an electron(s) to a nonmetal.

• The metal then becomes a(n) ____________ (positive ion). and the nonmetal becomes a(n) _____________ (negative ion).

• Their opposite charges cause them to combine into a crystalline solid.

PRACTICE: Determine the molecular formula of the compound formed from each of the following ions.

a. K+ & P3-

b. Sn4+ & O2-

c. Al3+ & CO32-



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CONCEPT: NAMING MOLECULAR COMPOUNDS Features: _________________ & _________________

Because molecular compounds combine in different proportions to form different compounds, we must use numerical

prefixes.

Rules for Naming: a. The first nonmetal is named normally and uses all numerical prefixes except ___________________. b. The second nonmetal keeps its base name but has its ending changed to _____________________. EXAMPLE: Write the formula for each of the following compounds.

a. Disulfur monobromide b. Iodine Tetrachloride

PRACTICE: Give the systematic name for each of the following compounds:

a. CO b. N2S4 c. IO5



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CONCEPT: IONIC COMPOUNDS

In the early days of chemistry, newly discovered compounds were given fancy names such as morphine, quicklime and

muriatic acid. Since then thousands of new compounds have been discovered and named under a system called

_____________________________.

Metals tend to __________ electrons to become positively charged ions called _______________.

Nonmetals tend to __________ electrons to become negatively charged ions called _______________.



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CONCEPT: NAMING BINARY IONIC COMPOUNDS Features: ___________________ & ___________________

Rules for Naming: a. The metal is named and written first.

• If the metal is a transition metal we must use a _________________________ to describe its positive charge.

b. The nonmetal keeps its base name but has its ending changed to ___________________.

EXAMPLE: Provide the molecular formula or name for each of the following compounds.

a. Calcium phosphide b. CoO

PRACTICE: Provide the molecular formula or name for each of the following compounds.

a. AlBr3 b. Lead (IV) sulfide c. SnO2



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CONCEPT: NAMING IONIC COMPOUNDS w/ POLYATOMICS Features: _________________ & _________________

Rules for Naming:

a) The metal keeps its name and is named and written first.

• If the metal is a transition metal we must use a _____________________ to describe its positive charge.

b) Name the polyatomic as you would normally.

EXAMPLE: Write the formula for each of the following compounds:

a. Iron (III) Acetate b. Copper (I) phosphate

c. Strontium Carbonate d. Ammonium Nitrite

EXAMPLE: Give the systematic name for each of the following compounds:

a. Pb(CrO4)2 b. Ga(ClO4)3

c. Mn(HSO4)2 d. Ba(CN)2



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CONCEPT: NAMING IONIC HYDRATES Features: _________________ & _________________

CuSO4 5 H2O

Rules for Naming the Ionic Compound portion: a. The metal is named normally and written first.

• If the metal is a transition metal we must use a ________________________ to describe its positive charge.

b. The nonmetal keeps the first part of its name but has its ending changed to ___________________.

c. Name the polyatomic as you would normally.

Rules for Naming the H2O portion:

a. The H2O portion will be called ___________________ .

b. To describe the number of H2O molecules use these prefixes.

EXAMPLE: Write the formula for each of the following compounds.

a. Calcium carbonate hexahydrate

b. Lead (IV) Sulfate pentahydrate


a) K2Cr2O7 · 3 H2O b) Sn(SO3)2 · 4 H2O



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CONCEPT: NAMING ACIDS

1. BINARY ACIDS Features: _______________________ + _______________________

Rules for Naming:

a. The prefix will be ___________________ .

b. Use the base name of the nonmetal.

c. The suffix will be ___________________ .

EXAMPLE: Write the formula for each of the following compounds:

a. Hydroiodic acid b. Hydroselenic acid c. Hydrofluoric acid


a. HBr b. H2S c. HCN

2. OXOACIDS or OXYACIDS Features: _______________________ + _______________________

Rules for Naming: a. If the polyatomic ion ends with –ate then change the ending to _____________________. b. If the polyatomic ion ends with –ite then change the ending to ______________________.

EXAMPLE: Give the systematic name or formula for each of the following compounds:

a. H2CO3 b. Nitric acid c. H2SO4

PRACTICE: Give the systematic name or formula for each of the following compounds:

a. Hypobromous acid b. HClO3 c. Acetic acid



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CONCEPT: COMMON POLYATOMIC IONS

Polyatomic ions are compounds made up of different elements, usually only ____________, and possess a ____________.

Singly Charged Cation (Positive Ion)

NH4+

Ammonium

Doubly Charged Cation (Positive Ion)

Hg22+

Mercury (I)

Singly Charged Anions (Negative Ions)

CH3CO2– or C2H3O2

–

Acetate

CN–

Cyanide

OH–

Hydroxide

MnO4–

Permanganate

NO3–

Nitrate

Nitrite

Doubly & Singly Charged Anions (Negative Ions)

HPO42–

Hydrogen Phosphate

H2PO4–

Dihydrogen Phosphate

HCO3–

Hydrogen Carbonate or Bicarbonate

HSO4–

Hydrogen Sulfate or Bisulfate

Doubly Charged Anions (Negative Ions)

CO32–

Carbonate

CrO42–

Chromate

Cr2O72–

Dichromate

O22–

Peroxide

SO42–

Sulfate

Sulfite

Triply Charged Anions (Negative Ions)

PO43–

Phosphate

Phosphite



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CONCEPT: POLYATOMIC IONS w/ HALOGENS

Polyatomic ions containing halogens are sometimes referred to as __________ halogens or halogen ________________.

These compounds share 4 common characteristics:

1.

2.

3.

4.

These compounds use the same system for naming:

PRACTICE: Name each of the following compounds.

a. BrO4 – b. FO2 –

c. ClO – d. IO3 –



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CONCEPT: DIPOLE ARROWS

Before drawing covalent compounds we first need to understand the idea of polarity and its connection to electronegativity.

• Polarity arises whenever two elements are connected to each other and there is a significant difference in their

electronegativities.

• Generally, electronegativity ________________ going from left to right of a period and ________________ going

down a group.

To show this difference in electronegativity we use a dipole arrow.

The dipole arrow points towards the ________________ electronegative element.

The Effect of Electronegativity Difference on Bond Classification

Electronegativity Difference (ΔEN)

Bond Classification

Example

Zero (0.0)

Pure Covalent

Small (0.1 – 0.4)

Nonpolar Covalent

Intermediate (0.4 – 1.7)

Polar Covalent

Large (Greater than 1.7)

Ionic



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PRACTICE: DIPOLE ARROWS

EXAMPLE: Based on each of the given bonds determine the direction of the dipole arrow and the polarity that may arise.

a. H Cl

b. S O

c. Br B Br

PRACTICE 1: Based on the given bond determine the direction of the dipole arrow and the polarity that may arise.

a. H C


a. N F


a. H N H



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CONCEPT: CHEMICAL BOND IDENTIFICATION

PRACTICE: Answer each of the following questions dealing with the following compounds.

KBr NH3 F2 CaO NaClO

a. Which of the following compound(s) contains a polar covalent bond?

b. Which of the following compound(s) contains a pure covalent bond?

c. Which of the following compound(s) contains a polar ionic bond?

d. Which of the following compound(s) contains both a polar ionic bond and a polar covalent bond?



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CONCEPT: ENERGY CONSIDERATIONS IN IONIC BONDING

________________________ is the enthalpy change that occurs when 1 mol of ionic solid separates into gaseous ions. It

tells us the strength of ionic interactions and has an influence in melting point, hardness, solubility and other properties.

Li+ (g) + F – (g) LiF (s) ∆H = 1050 kJ/mol

In order to calculate the energy of an ionic bond we use the following equation;

Ionic Bond Energy =

Radius = __________________________________

EXAMPLE: For each pair, choose the compound with the lower lattice energy.

a. BaO or MgO

b. LiCl or CaS

PRACTICE 1: Choose the compound with the lower lattice energy.

a. AlN or KBr

PRACTICE 2: Choose the compound with the higher lattice energy.

a. CsF or LiCl



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CONCEPT: LATTICE ENERGY APPLICATION

Lattice Energy represents the energy released when 1 mole of an ionic crystal is formed from its gaseous ions.

Lattice Energy (Electrostatic Energy) =Cation Charge ⋅Anion ChargeCation Radius +Anion Radius

Mg2+(g) + O2−(g)→MgO (s) ΔH = −3800 kJmole

EXAMPLE 1: The solubilities of CaCrO4 and PbCrO4 in water at 25°C are approximately 0.111 g/L H 2O and 0.0905 g/L H2O

respectively. Based on this information, which compound do you think has the smaller lattice energy?

EXAMPLE 2: Which of the following bond will have the highest ionic character?

A. BeBr2 B. MgBr2 C. SrBr2 D. BaBr2

Increases

Increases

Lattice EnergyMg

O

_________ ion charges and _________ radii help to increase lattice energy.

The larger the lattice energy then the stronger the ionic bond between the ions.

Results in a higher boiling point and melting point for the ionic compound.

Generally, it increases going from left to right of a period and increases going up any group because of a(n) _________ in atomic size.



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CONCEPT: BORN-HABER CYCLE

The Born-Haber cycle is used a method to calculate the ________________________ or _______________________ of a compound.

• It looks mainly at the formation of an ionic compound from gaseous ions.

• The metal being from Groups _______ or _______ and the nonmetallic element being a ________________ or ________________ .

M (s)+ 12X2

ΔHfo

⎯ →⎯⎯ MX (s)

M (g) X (g)

M+ (g) X– (g) HX (s)

1

2

3

4

5

ΔHfo = 1 2 3 4 5+ +++

1

2

3

4

5

=

=

=

=

=

+



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PRACTICE: BORN-HABER CYCLE

EXAMPLE: Using the Born-Haber Cycle, demonstrate the formation of cesium chloride, CsCl, and calculate its heat of formation.

ΔHSublimation = 79kJmol

IE1 = 376kJmol

ΔHDissociation =122kJmol

EA = −349 kJmol

U = −661 kJmol



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CONCEPT: ELECTRON-DOT SYMBOLS

Before we look at the first two bonding models, we have to figure out how to depict the valence electrons of bonding atoms.

• In the _________ electron-dot symbol, the element symbol represents the nucleus and inner electrons, and the

surrounding dots represent the ________________ electrons.

EXAMPLE: Draw the electron-dot symbol for each of the following elements.

1A 2A 3A 4A 5A 6A 7A 8A

Li

Be

B

C

N

O

F

Ne

It’s easy to write the Lewis symbol for any Main-Group element:

1) Remember that Group Number equals Valence Electron Number.

2) Place one dot at a time on the four sides (top, right, bottom, left) of the element symbol.

3) Keep adding dots, pairing them up until you have reach the number of total valence electrons for that element.

PRACTICE 1: Draw the electron-dot symbol for the following ion.

Mg2+


N3-


Cr1+



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CONCEPT: CHEMICAL BONDING I

Rules for Drawing

1. Least electronegative element goes into the center. Important Facts to Know:

(a) Electronegativity increases across any Period going from left to right and up any Group going from bottom to top.

(b) Hydrogen and Fluorine ________________ go in the center and they only make _________ BOND.

2. Number of valence electrons equals group number.

3. Carbon must make _____ bonds, except in rare occasions when it makes _____ bonds.

• If the carbon atom were positive or negative then it would make _____ bonds

4. Nitrogen likes to make _____ bonds.

5. Oxygen likes to make _____ bonds.

6. Halogens (Group 7A), when not in the center, make _____ bond.

7. Expanded Valence Shell Theory: Nonmetals starting from Period _____ to _____ can have more than 8 valence

electrons around them when in the center.



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CONCEPT: INCOMPLETE OCTETS

Nonmetals form covalent bonds to generally follow the ___________ rule, in which the element is surrounded by 8 valence

electrons.

• Sometimes elements form compounds in which they have ____________________ 8 valence electrons.

• These elements are said to have an incomplete octet or to be ________________________________________ .

EXAMPLE: Draw the following molecular compound.

BH3

PRACTICE: Draw the following molecular compound.

BeCl2



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CONCEPT: EXPANDED OCTETS

Expanded Valence Shell Theory: Nonmetals starting from Period _____ to _____ can have more than 8 valence

electrons around them when in the center.

EXAMPLE: Draw each of the following molecular compounds.

IF3 KrF5+

PRACTICE 1: Draw the following molecular compound.

SBr4


I3–



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CONCEPT: POLYATOMIC IONS

Shortcut: If you have _____, _____, _____, _____, __________________ or __________________ connected to oxygen

then the negative charge tells us how many oxygens are single bonded.

• The remaining oxygens are _______________________ bonded to the central element.

EXAMPLE: Draw each of the following molecular compounds.

SO42-

PO43- H2SO4


SeO42-


XeO64-



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CONCEPT: FORMAL CHARGE

Structures and polyatomic ions that break the octet rule often have ________________ Lewis Structures.

• The purpose of using the formal charge formula is to determine which Lewis structure is the best answer.

Formal Charge =

a) Use formal charge formula to check to see if you drew your compound correctly.

b) Formal charges must be either _____, ______, ______.

c) If you add up all the formula charges in your compound that will equal the overall charge of the compound.

EXAMPLE: Calculate the formal charge for each of the following element designated for each of the following.

a. The carbon atom in

b. The sulfur atom in

PRACTICE: Calculate the formal charge for each of the following element designated in the following compound.

a. Both oxygen atoms in:

!A B



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CONCEPT: RESONANCE STRUCTURES

Resonance structures are used to represent bonding in a molecule or ion when a single Lewis structure cannot correctly

describe the Lewis structure.

EXAMPLE: Determine all the possible Lewis structures possible for NO2–. Determine its resonance hybrid.

EXAMPLE: Determine the remaining resonance structures possible for the following compound, CO32-.

O

C OO



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16. How many milligrams of NaCN are required to prepare 712 mL of 0.250 M NaCN?



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17. What volume (in µL) of 0.100 M HBr contains 0.170 moles of HBr?



Page 28

18. How many moles of Ca2+ ions are in 0.100 L of a 0.450 M solution of Ca3(PO4)2?



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19. How many chloride ions are present in 65.5 mL of 0.210 M AlCl3 solution? a) 4.02 × 1023 chloride ions

b) 5.79 × 1024 chloride ions

c) 2.48 × 1022 chloride ions

d) 8.28 × 1021 chloride ions

e) 1.21 × 1022 chloride ions



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22. To what final volume would 100 mL of 5.0 M KCl have to be diluted in order to make a solution that is 0.54 M KCl?



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23. If 880 mL of water is added to 125.0 mL of a 0.770 M HBrO4 solution what is the resulting molarity?



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26. Consider the following balanced redox equation:

H2O + 2 MnO4 – + 3 SO32- 2 MnO2 + 3 SO42- + 2 OH – How many grams of MnO2 (MW: 86.94 g/mol) are produced when 32.0 mL of 0.615 M MnO4- (MW: 118.90 g/mol) reacts with excess water and sulfite?



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27. Iron (III) can be oxidized by an acidic K2Cr2O7 solution according to the net ionic equation:

Cr2O72- + 6 Fe2+ + 14 H+ 2 Cr3+ + 6 Fe3+ + 7 H2O

If it takes 35.0 mL of 0.250 M FeCl2 to titrate 50 mL of a solution containing Cr2O72-, what

is the molar concentration of Cr2O72-?



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28. Vinegar is a solution of acetic acid, CH3COOH, dissolved in water. A 5.54 g sample of

vinegar was neutralized by 30.10 mL of 0.100 M NaOH. What is the percent by weight of

acetic acid in the vinegar?



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29. What is the molar mass of a 0.350 g sample of a monoprotic acid if it requires 50.0 mL of 0.440 M Ca(OH)2 to completely neutralize it?



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30. Give the complete ionic equation for the reaction (if any) that occurs when aqueous solutions of sodium sulfide and copper (II) nitrate are mixed. a) Na+ (aq) + SO42-(aq) + Cu+(aq) + NO3-(aq) CuS(s) + Na+(aq) + NO3-(aq)

b) Na+ (aq) + S-(aq) + Cu+(aq) + NO3-(aq) CuS(s) + NaNO3(aq)

c) 2 Na+(aq) + S2-(aq) + Cu2+(aq) + 2 NO3-(aq) Cu2+(aq) + S2-(aq) + 2 NaNO3(s)

d) 2 Na+(aq) + S2-(aq) + Cu2+(aq) + 2 NO3-(aq) CuS(s) + 2 Na+(aq) + 2 NO3-(aq)

e) No reaction occurs.



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31. Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of H2SO4 and KOH are mixed. a) H+(aq) + OH-(aq) H2O(l)

b) 2 K+(aq) + SO42-(aq) K2SO4(s)

c) H+(aq) + OH-(aq) + 2 K+(aq) + SO42-(aq) H2O(l) + K2SO4(s)

d) H22+(aq) + OH-(aq) H2(OH)2(l)




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32. Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of Na2CO3 and HCl are mixed. a) 2 H+(aq) + CO32-(aq) H2CO3(s)

b) 2 Na+(aq) + CO32-(aq) + 2 H+(aq) + 2 Cl-(aq) H2CO3(s) + 2 NaCl(aq)

c) 2 H+(aq) + CO32-(aq) H2O(l) + CO2(g)

d) 2 Na+(aq) + CO32-(aq) + 2 H+(aq) + 2 Cl-(aq) H2CO3(s) + 2 Na+(aq) + 2 Cl-(aq)




Page 39

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CH.4 - CHEMICAL BONDING: UNDERSTANDING CLIMATE CHANGElightcat-files.s3.amazonaws.com/packets/admin... · • Generally, the reason is that ionic bonding _____ the potential energy