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AU / Chemistry / CHEM 1030 / Are ionic compounds the same as ionic bonds?

Are ionic compounds the same as ionic bonds?

Are ionic compounds the same as ionic bonds?

Description

School: Auburn University
Department: Chemistry
Course: Fundamentals Chemistry I
Professor: John gorden
Term: Fall 2015
Tags:
Cost: 50
Name: CHEM 1030
Description: This is the study guide for Exam @ in CHEM 1030.
Uploaded: 10/25/2016
9 Pages 68 Views 5 Unlocks
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Exam 2 Study Guide Chapters 5 - 7 CHEM1030


Are ionic compounds the same as ionic bonds?



Since the tests seem to be heavily concept-based, I've compiled some concepts at  the end of every chapter review that might be important to understand.

∙ Chapter 5

o Ionic Compounds and Ionic Bonding

 Formed by ionization (complete transfer of electron from one  atom to another atom)

 Ionization causes ions to form a crystalin structure called a  

lattice

∙ Lattice energy is the energy it takes to convert one  We also discuss several other topics like What is a price maker?

mole of a solid ionic compound into a gas.

 When naming anions (negatively charged atoms), add the  

ending "-ide" to the element name.

∙ Ex. Cl- is chloride


How do you calculate molecular mass and formula mass?



 When naming cations (negatively charged particles), add the  word "ion" after the element name

∙ Ex. Na+ is sodium ion

 Remember when making ionic compounds, the resulting charge  must be neutral.

∙ Ex. Write the formula for magnesium fluoride

Magnesium has +2 charge and fluoride has -1  

charge.

Make the charges equal zero to make a neutral  

coompound.

(2) + (-1)(2) = 0

As seen by the formula, fluorine needs a +2 charge to  Don't forget about the age old question of Why use bacteria and bacteriophage model systems?

make the compound neutral

MgF2

o Polyatomic Ions

 Ions that have more than one charged atom bonded together.  Below is a chart of polyatomic anions to memorize

Common Polyatomic Ions

Name

Formula/Charge

Cations

ammonium

NH4+

hydronium

H3O+

mercury (I)

Hg22+

Anions

acetate

C2H3O2-

azide

N3-

carbonate

CO32-

chlorate

ClO3-


How do you calculate lattice energy?



If you want to learn more check out Which famous italian family were huge patrons of the arts during the renaissance?

Exam 2 Study Guide Chapters 5 - 7 CHEM1030

chromate

CrO42-

cyanide

CN 

dichromate

Cr2O7-

dihydrogen phosphate

H2PO4-

hydrogen carbonate of bicarbonate

HCO3-

hydrogen phosphate

HPO22-

hydrogen sulfate

HSO4-

hydroxide

OH 

hypochlorite

ClO 

nitrate

NO3-

nitrite

NO2-

oxalate

C2O42-

perchlorate

ClO4-

permanganate

MnO4-

peroxide

O22-

phosphate

PO43-

phosphite

PO33-

sulfate

SO42-

sulfite

SO32-

thiocyanate

SCN-

If you want to learn more check out What is residence time?

 Rules to help with memorization (DON"T FORGET TO MEMORIZE  CHARGES):

1. The "-ate" form of the ion has one more oxygen atom  

than the "-ite" form.

2. The "hyper – ite" form of the ion has one less oxygen than

the "-ite" form. Don't forget about the age old question of Where do good ideas come from?

3. The "per – ate" form of the ion has one more charge than  

the "-ate" form.

 When asked to balance an ionic compound, you must balance  the charges. If you want to learn more check out What is true about communication in men and women??

 Use the stock system (uses roman numerals to indicate  

charge) if the metal is a trasition metal ('d' orbital element)

o Organic Compounds

∙ Compunds containing a carbon atom

∙ Name of compound depends on number of carbon atoms in molecule  Below is a chart of the alkanes (simple hydrocarbons) you must memorize

 There is a chart on page 163 of organic compounds you need to  memorize.

∙ A friend told me they memorize the compounds by  

thinking "Me Eat Peanut Butter" for the first 4 (methane,  

ethane, propane, butane). After those, the next 6 use the  

common prefix names that tell how many carbon atoms  

there are with the name "-ane" added to the end of the

Exam 2 Study Guide Chapters 5 - 7 CHEM1030

word (pentane, hexane, heptane, octane, nonane,  

decane)

∙ Also, the formula for knowing the formula of an organic  

compound is  

C4H(2n + 2)

o Ex. If asked to write the molecular formula for  

butane:

There are 4 carbons, and if you plug 4 into the  

subscript  

formula you get 10 hydrogens.

SO

C4H10 is the formula for butane.

∙ Molecular Formulas and Formula Masses

 To find molecular mass of a compound, add the molar masses  of each elemental ratio together.

Molecular Mass = Σ(number of atoms of element) x  

(atomic mass of element)

∙ Ex. Find molecular mass of H2O.

MM = 2(1.001) + 15.999 = 18.001 g

 To find percent composition of one element in a compound,  use a ratio of the elemental mass and the total mass

% Composition = ((mass of element) / (mass of total  

compound)) x 100

∙ Ex. Find the percent composition of oxygen in a water  

molecule (H2O).

% O = (15.999) / ((15.999) + (2 x 1.001)) = 5.926%

 To determine the empirical formula from the percent  

composition:

1. Find grams of each element

2. Convert grams of each element to moles

3. Divide the molar amounts of by the smallest molar  

amount o all elements involved (this will give you the  

whole number ratio of the empirical formula

∙ Ex. Determine the empirical formula of a compound that  

is 30.45% N and 69.55% O.

First, use the percent composition as the masses of

each element (we

Exam 2 Study Guide Chapters 5 - 7 CHEM1030

can do this if we assume the molecular mass is  

100g).

30.45% N = 30.45 g N

69.55% O = 69.55 g O

Now, convert molar mass to moles.

30.45 g N x (1 mol / 14.009 g N) = 2.173 mol N

69.55 g O x 1 mol/15.999 g O = 4.347 mol O

Now divide each molar amounr by the smallest of  

the molar amounts.

2.173/2.173 = 1 mol

4.347/2.173 = 2.0004 = 2 mol

Write the completed epirical formula with whole  

number subscripts.

Empirical Formula: NO2

∙ If you are given the molar mass of the compound, you can

then find the molecular formula (which includes  

subscripot ratios of elements as they actually appear).  

There wasn't an example of this in class, but it was on  

ALEKS and in the book, so I'll put it on the study guide in  

case.

So, say the question tells you that the molar mass of the  

compound is 92 g/mol.

Divide the empirial formula's mass by the  

molecular formula's mass to get a whole number.

Empirical Formula Mass = 14.001 + (2 x 15.999) = 46.01.

92/46.01 = about 2.

Multiply each subscript in the mpirical formula by  

thiss whole number ratio.

Molecular Formula: N1 x 2O2 x 2 = N2O4

o Concepts From Chapter 5

 When ionization occurs, the process of pulling an electron away from an atom requires energy.

∙ Ionization energy increases with increasing  

electronegativity.

Exam 2 Study Guide Chapters 5 - 7 CHEM1030

1. Lattice energy is a measure of compounds stability. So the  higher the lattice energy, the higher the stability of the ionic  

compound.

∙ Lattice energy is proportional to the distance between  

atoms. The shorter the distance between atoms, the  

highter the lattice energy.

Exam 2 Study Guide Chapters 5 - 7 CHEM1030

∙ Chapter 6

o Lewis Dot Structure Theory says atoms will lose or gain electrons in order to obtain a noble gas configruation (more stablility).

o Rules for drawing Lewis Structures for 2nd period elements: 1. Count the number of total number of valence electrons in the  compound.

2. Place double bonds and surrounding elements around the  

central atom of a molecule (the central atom is usually the least  electronegative atom UNLESS hydrogen is involved).

3. Subtract the number of electrons you've already drawn from the  total number of valence electrons. This difference tells how  

many electrons you still need in your structure.

4. Place remaining electrons around the surrounding elements.  Make sure to add double bonds where needed to make atoms  

fuflill the octet rule.  

Atoms in the 3s orbital and higher can hold more than 8  

electrons because the 'd' orbitals can hold electrons.

o If the compound contains an added postive charge, make sure you  accommodate one (or more) less electron into your structure. If the  compound contains an added negative charge, make sure you  

accommodate one (or more) more electron into the structure.

 If the compound contains any added charge, make sure you  

indicate this with brackets around the structure with the added  

charge.

o Sometimes there are multiple ways to draw the same Lewis Structure.  This results in resonance structures.

o Concepts From Chapter 6

 Atoms lose or gain electrons to gain a noble gas configuration  for more stability.

 Resonance structures make a molecule more stable. So, the  

more resonance structures there are, the more stable a  

molecule is.

∙ Chapter 7

o Valence-Shell Electron-Pair Repulsion (VSEPR) Model of atomic  structure uses the fact that electron domains will get as afar away from each other as possible to get into a more stable structure.

1. Electron Domain Geometries  

∙ Geometry depends on number of electron domains  

(including lone paris of electrons)

∙ 5 Electron Domain Geometries

1. Linear - 180°

2. Trigonal Planar - 120°

3. Tetrahedral – 109.5°

Exam 2 Study Guide Chapters 5 - 7 CHEM1030

4. Trigonal Bipyrimidal - 90°, 120°

5. Octahedral - 90°

 Molecular Geometries

∙ Geometry depends on number of electron domains AND if  

lone pairs of electrons or double bonds are present.

∙ The chart on page 226 in the book lists all of the  

molecular geometries based on the electron domain  

geometries. (Dr. Blumenthal didn't go in depth about any  

of these. Actually, he only went over ONE of them, so I  

wasn't sure if he expected us to know all of them or not.  

Just in case, I have at least mentioned them).

∙ Molecular geometries that differ from electron domain  

geometries are caused by deviation from ideal bond  

angles.

o Lone pair of electrons and double bonds between  

atoms repel surrounding atoms more effectively.

This causes angles that differ slightly from the  

previously listed ideal angles.

o Polarity results from an unequal sharing of electrons due to a  difference in electronegativity or atoms.

 Polarity occurs when a dipole moment (when an atom has a  partial negative or positive charge due to electronegativity)  

exists

∙ Atoms that are more electronegative are more negative  

(because they are more attracted to electrons).

∙ Rules for determining polarity of a molecule:

1. Draw the Lewis Dot structure

2. Determine which atoms will have a partial negative  

charge and which ones will have a partial positive  

charge.

3. Draw corresponding arrow directions. On a graph.

4. Use vector sum rules to add the vectors (this is  

hard to describe in words, but I drew it on my  

uploaded chapter 7 notes).

Exam 2 Study Guide Chapters 5 - 7 CHEM1030

5. If overall vector sum is zero, there is NO dipole  

moment, ehich means the moleculs is nonpolar.

If the overall vector sum is NOT zero, there IS a  

dipole moment, which means the molecule is polar.

o Valence Bond Theory uses quantum mechanic description of orbitals (meaning it uses the actual shapes of different orbitals) to show a 3-D  structure of a molecule.

 Utilizes fact overlapping orbitals can hold single electron to bond  Hybridization is the mixing of atomic orbitals to show how  

atoms can make double bonds. (It is a continuation of the  

valence bond theory).

 It is almost impossible for me to show this process in depth  

without pictures, although it is in the uploaded written chapter 7 notes.  

However, here is an overview of the rules:

1. Draw the Lewis Dot Structure of the given molecule

2. Observe the geometry using the VSEPR model. The  

number of electron domains in the Lewis structure is the  

number of hybrid orbitals you should end up with.

3. Write out the electron configuration of the central atom. If

there are any paired electrons, move one of the electrons  

in each pair to a higher orbital.

4. Combine the orbitals as necessary.

5. Once hybridized central atom is drawn, add the  

surrounding atoms in their respective orbital shape.

Once the structure is drawn, it will contain sigma bonds, bonds that occur where probability of electron-sharing is high (usually  

directly between nuclei), and pi bonds, bonds that occur where  there is a low probability of electron-sharing (not directly  

between nuclei).

 Molecular Orbital Theory uses diagrams to describe if a  

diatomic molecule is paradiameagnetic (has unpaired  

electrons) or diamagnetic (has all paired electrons).

∙ States that orbitals combine to make NEW orbitals called  

bonding molecular orbitals and antibonding  

molecular orbitals

∙ When filling out the diagram, remember to follow Hund's  

Rule

Exam 2 Study Guide Chapters 5 - 7 CHEM1030

∙ Bond Order can be determined from the resulting  

diagram

BO = ½(# of electrons in bonding molecular  

orbitals – # of electrons in antibonding molecular  

orbitals)

o Concepts From Chapter 7

 VSEPR Model: uses number of electron domains around central atom to determine structure and angles of structure. Provides  

relative structure of molecule.

 Valence Bond Theory: explains why covalent bonds form: to  reduce the overall potential energy of the isolated atoms (in  

other words, to become more stable)

 Hybridization: continuation of valence bond theory. Adds  

explanation of why double bonds can form in molecules (double  bonds form when one sigma bond is combined with one pi  

bond).

 Molecular Orbital Theory: uses diagrams that can describe  important properties of diatomic molecules (paramagnetism  

and diamagnetism).

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