Study Guide Chemistry Comprehensive Exam 2014 I

Study Guide Chemistry Comprehensive Exam 2014

1 Made By Ashley Thomas

I. Memorization: see “Stuff to Memorize Sheet” for more (from teacher)

Ions that Form Soluble Compounds

Group 1 ions (Li+, Na+, ect.)

Ammonium

Nitrate

Acetate

Bicarbonate (aka Hydrogen Carbonate)

Chlorate

Perchlorate

Polyatomic ions

Nitrate

Chlorite

Hydrochlorite Chromate

Phosphite

Dichromate

Cyanide Chlorate

Hydroxide Sulfate

carbonate

Bicarbonate*

Sulfite

Permanganate

Phosphate

*** Acetate

Perchlorate

Ammonium

*AKA Hydrogen Carbonate

II. Lab stuff

Accurate: A measurement that is close to the accepted or correct value.

Precise: Measurements that are close to one another.

Percent Error: Measures accuracy.

Uncertainty: Measures precision.

Absolute Deviation: The difference from the average. Report using

Percent Error:

Uncertainty: ( means average)

III. Significant figures

Types of Zeros

o Sandwich zero: 505

o Leading zero: 0.0005



o Trailing zero: 5000

Rules

o All non-zero digits are significant

o All sandwich zeros are significant

o Leading zeros are never significant

o Trailing zeros are not significant when there is no decimal and are always significant when

there is a decimal after of in-between.

Addition and Subtraction: You must round your answer to the least number of decimal places.

Multiplication and Division: You must round you answer to the least number of significant

figures.

IV. Physical and chemical changes/properties

Precipitate: A solid formed when two solutions are mixed.

Spectator Ion: An ion that doesn’t participate in a chemical change.

Physical Change: In a physical change molecules stay together, the substance doesn’t change,

bonds are never broken, and the change is reversible.

Chemical Change: In a chemical change bonds are broken and reformed, new substances are

formed, and the change may not be reversible.

V. Periodic table basics

Groups

o 1a/b-Alkali Metals

o 2a/b- Alkaline Earth Metals

o 3b-12b- Transition Metals

o 7a/17b- The Halogen

o 8a/18b- Noble Gasses

o FOR CHARGES SEE PERIODIC TABLE 1

Metals vs. Metalloids vs. Non-metals

o Metals: to the left of the stairstep

o Non-metals: to the right of the stairstep

o Metalloids: Directly touching the stairstep

Ionic compound: Metal and non-metal

Covalent Compound: only non-metals

VI. Periodic table trends and properties (to compare and explain)

Why are orbital diagrams so important? They allow you to predict the ions of elements,

including transition metals.



Valence Electrons: The outermost electrons in an atom, that are included in chemical reactions.

Core Electrons: The inner electrons of an atom, that aren’t included in chemical reactions.

Another name for core electrons is shielding electrons.

o Shielding electrons “shield” the outermost electrons from the nucleus “pull”.

o Called effective nuclear charge ( )

Z is the atomic number or number of protons

S is the number of shielding electrons

o The higher the the more the valence electrons are attracted to the nucleus.

o The higher the the more likely the element will become an anion(negative ion).

Periodic Trends (See Periodic Table 2)

o Homework

Why does atomic radius decrease as you move across the periodic table?

As you move from right to left is increasing. The nucleus “pulls” electrons in

closer as a result of the higher .

How are the trends of atomic radius and ionization energy related?

They are inversely related, so as ionization energy increases, atomic radius

decreases, left to right in a period. In a group, as you go down, ionization energy

decreases and atomic radius increase.

What’s the difference between electronegativity and ionization energy?

Ionization energy removes electrons and electronegativity is abou the attraction of

electrons.

What does it mean to be isoelectronic? Identify the largest ion and smallest ion of the

group: . Explain why.

Isoelectronic means same number of electrons.

Ex) They all have 10 electrons.

Smallest Ion: .

Largest Ion:

The number of protons is decreasing for the same number of electrons.

The first ionization energy of beryllium is 9.322 eV, the second ionization energy is

18.211 eV, and the third ionization energy is 153.893 eV. Explain why the third

ionization energy of beryllium is so much higher than the first two.

The first two are valence electrons and the next is a core electron.

o Atomic Radius/size: see above notes

o Ionic Radius/size: The radius of an ion.

For cations: the ion of an atom is smaller than the atom. Reason: same number of

protons, fewer electrons, so more “pull”.

For anions: the ion of an atom is larger than the atom. Reason: same number of

protons, more electrons, so less “pull”.

o Electronegativity/Electron Affinity: How much an atom attracts an electron.

o Ionization Energy (1st): The energy required to remove an electron from an atom.



1st IE: The energy required to remove the 1st or outermost electron.

o Metallicity/Metallic Character: How much like a metal an element is.

o Reactivity: The ability to gain or lose electrons more easily.

Metals reactivity increase as you go left and down a group.

Reactive Metals: easily lose their electrons; low ionization energy; large atomic

radius; low electronegativity.

Low ionization energy: most important characteristic.

Non-metals reactivity increases as you go right and up a group.

Reactive Non-metals: easily gain electrons; high ionization energy; small atomic

radius; high electronegativity.

High electronegativity: most important characteristic.

Factors affecting periodic trends:

o : effective nuclear charge (only explains trends across a period); every atom in the same

group has the same

o Distance/Size: The farther from the nucleus, the less “pull” the nucleus can exert. (Columb’s

Potential Energy)

o Columb’s Potential Energy:

k- Constant -Charges d- distance V- potential energy

Takes into account both charge and distance.

According to Columb’s PE, as distance decreases, potential energy will increase. If

potential energy increases then the electron is more attracted to the nucleus.

If q (charge o the nucleus) increases the potential energy increases.

VII. Naming Compounds

Charges:

o Zn (zinc) is always 2+

o Ag (silver) is always 1+

Steps

1) Identify the type of compound. Then follow specific steps for each type.

Ionic: metal and non-metal (cation and anion); give and take electron. TIP: Does it have

a metal?

Covalent: non-metals only; shared electrons. TIP: Not allowed to have a metal, no H in

front.

Acid: TIP: H in front, no metals.

Covalent: prefix (no mono-) and name of 1st non-metal + prefix and name of second metal plus –

ide.

o Prefixes:

1 Mono 6 Hexa

2 Di 7 Hepta

3 Tri 8 Octa

4 Tetra 9 Nona

5 Penta 10 Deca



Ionic:

1) Binary or polyatomic?

Binary: 2 elements only.

Polyatomic: 3+ elements.

o Binary: Name of metal + name of non-metal plus –ide.

o Polyatomic: name of metal + mane of polyatomic ion.

o Transition Metal Compounds: Need a roman numeral in the name except for zinc (Zn) and

Silver (Ag). The Roman numeral represents the charge of the metal. Transition metals have

variable charges so they need a roman numeral.

Acidic: H in front, because acids donate H+ ions.

o HCl: Chloride hydro-chlor-ic acid

Hydro-name of anion- ic +acid

TIP: Hydro- only H and 1 element

o : Chlorite Chlor-ous acid

Name of anion- + acid

o :Chlorate Chlor-ic acid

Name of anion-ic + acid

o If you have a polyatomic ion in your acid, the acid’s name doesn’t include hydro.

VIII. Writing Chemical Formulas

: represents the atoms in a molecule (structure); subscripts tell us how many of each atom

is present; symbols tell us which elements are in the molecule.

Covalent:

1) Write the symbol of the elements

2) Write the subscripts of the elements according to the prefixes

Ionic:

1) Write the symbol of the cation and its charge.

2) Write the symbol of the anion and its charge.

-ide element (except hydroxide and cyanide)

-ite, -ate anion is polyatomic ion

3) Change the subscripts so the charges add up to zero

o Memory Device: Symbol-Charge, Symbol-Charge, cross it, make it pretty

o TIPS:

Make sure the charges add to zero

Reduce subscripts if possible

Parentheses around polyatomic ions

Roman numeral is the charge not the subscript

Acidic: H in from all the time

1) Write the symbol of the anion with charge

2) Add enough H+ to make the charge add up to zero

o TIPS:



Use the suffix to determine the anion

Hydro+ic just one element

+ic polyatomic ion ending in –ate

+ous polyatomic ion ending in –ite

IX. Percent composition

Converting Moles (mol.) to grams (g):

Molar mass of a single element is the mass found on the periodic table

X. Empirical Formulas

Empirical Formula: Represents the ratio between the number of each atom found in the

compound. The ratio is a mole to mole ratio.

Ex) 1 mol C atoms: 1 mol H atoms: 3 moles of Cl atoms

Empirical formulas don’t always represent the actual structure of the compound. They only

tell you the correct ratio.

There are 3 possible starting points for an empirical formul calculation:

(1) % comp (2) Grams (3) Moles Mole Ratio (subscripts)

(1) (2) : change percent to grams (assume 100 grams)

(2) (3) : divide by molar mass

Step 4 (mole ratio): divide all by smallest number of moles

Why assume 100g? Makes conversion simple.

Why divide by the smallest number of moles? Allows you to make a mole ratio where the

smallest number is 1.

TIPS:

What if the ratio isn’t a whole number?

Ex) 1.5 mol O *2 = 3

1 mol N *2= 2

Ends in….

.1 or .9 5ound up/down

.2 or .8 multiply by 5



.5 multiply by 2

Example Problem: Calculate the empirical formula that contains 38.6% N and 63.3% O.

Molecular Compounds: Share the same ratio of atoms, but their actual structures differ. How

do we tell the difference between compounds?

Represents the actual structure of the compound but shares the same ratio as the

empirical formula.



Easy way to remember the difference between types of formulas:

Empirical: simplest ratio, not structure

Molecular: structure and ratio

o How to convert between empirical and molecular formulas:

Multiply subscripts of empirical by the whole number

o Example Problem: The compound ethylene glycol is used in antifreeze. IT contains 38.7% C,

9.75% H, and the rest is oxygen. The molecular weight of ethylene glycol is 62.07 g/mol.

What is the molecular formula?

XI. Balancing Chemical Equations

Conservation of Mass: Mass can’t be created or destroyed.

Conservation of mass tells you two important things about equations:

1) The amount of each element must be the same on both sides of the equation.

2) The types of elements must also be the same on both sides of the equation.

We balance reactions to satisfy the Law of Conservation of Mass.

What are you allowed to do to balance an equation?

o Change the coefficients of compounds

What are you not allowed to do to balance an equation?

o Change the subscripts of an element

o Change the elements

TIPS:

o Treat polyatomic ions as a unit or group

o Write water (H2O) as HOH

o No fractions in final answers

o Reduce coefficients only if you can reduce all of them

o For combustion reactions balance in the following order: H C O

XII. Chemical Equations/reactions

5 basic types:

1) Combustion

2) Synthesis

Oxidation- reduction

3) Decomposition


4) Single Replacement


5) Double Replacement

Acid-base



Precipitation

Oxidation Reduction (redox), acid-base (a/b), and precipitation (ppt) are special classes of the 5

main types

To predict the products of a reaction always identify the simple type (1 of the 5) first and then

worry about the special class.

Combustion (1): be able to write both reactants and products from words.

o A reaction where a compound, usually a hydrocarbon (mostly made of carbon), reacts with

oxygen gas (O2) to form carbon dioxide (CO2) and water (H2O)

Ex) CH4, methane, combusts. Write the reaction.

Synthesis (2): be able to identify the type; no prediction

o A reaction where simple reactants (elements or products) form a single compound.

o General form: A+B AB

Ex)

Decomposition (3): be able to identify the type; no prediction

o A reaction where one complex compound breaks into multiple simple products

o Opposite of synthesis

o General form: AB A+B

Ex)

Single Replacement/ displacement (4): be able to do simple product prediction (not if it occurs)

o General form: A + BC AC + B

Ex)

Ex)

o TIP: Cations switch with other cations, and anions switch with other anions.

o Driven by the reactivity of elements

Double Replacement/ displacement (5): be able to do complex product prediction and if the

reaction occurs.

A reaction where two ionic compounds switch their cations and anions

General form: AB +CD AD +CB

Ex) precipitate:

Ex) Acid-base:

Driven by the formation of a solid, liquid, or gas

Usually at least one of your reactants is a solution (aq: dissolved in water)

This allows compounds to ionize (separate into ions) and then switch their ions

Predicting products

o Double Replacement

1) Predict products

2) Check subscripts

3) Label States {TIP: Soluble (aq), insoluble (s)}

4) Balance (TIP: if all products are aq, there is no reaction)

o Combustion

is always the reactant

____ + +



XIII. Redox reaction

Oxidation-reduction (redox) have two processes that take place:

Oxidation: The loss of electrons. ex- Rust

Reduction: The gain of electrons. Ex- silver mirror, electroplating

These two processes always take place together

OIL RIG: Oxidation Is Lost, Reduction Is Gained

LEO the lion says GER: Lose Electrons is Oxidation, Gain Electrons is Reduction

Any reaction where electrons are transferred is a redox reaction

How so we tell when redox occurs? Oxidation number/state

The oxidation number of an atom/ion helps us keep track of electrons in a reaction.

Atoms/ions are oxidized or reduced, not compounds (An element not a compound)

Atoms/ions that become more positive are oxidized (Cu Cu2+)

Atoms/ions that become more negative are reduced (Cu2+ Cu)

o If oxidation numbers don’t change, there is no redox reaction.

o Oxidation Number/State Rules:

1) The oxidation number for an atom in its elemental form is always zero.

A substance is element if both of the following are true:

Only one kind of atom is present

There is no charge on the element

Subscripts don’t matter, the elements just have to be by themselves.

2) The oxidation number of a monoatomic (single type of atom) ion= charge of the

monatomic ion.

3) The oxidation number of all group 1 metals is 1+ (unless elemental).

4) The oxidation number of all group 2 metals is 2+ (unless elemental).

5) Hydrogen (H) has two possible oxidation numbers:

1+ when bonded to a nonmetal (covalent compounds)

1- when bonded to a metal (ionic compounds)

6) Oxygen (O) has two possible oxidation numbers:

2- in almost all compounds (peroxide is the exception)

1- in peroxides… very rare don’t need to worry about it

7) The oxidation number of fluorine (F) is always 1-.

8) The sum of the oxidation numbers of all atoms (or ions) in a neutral compound is 0.

9) The sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge

on the polyatomic ion.

Balancing

o Redox reactions must be balanced by both mass and charge because electrons are

transferred during the reaction.

o Balancing by charge conserves the amount of electrons transferred in the reacting

o To balance a redox reaction, we use the half reaction method

o Half Reaction Method: breaks the whole reaction into two parts

1) Oxidation half reaction: electrons go on the right (product)

2) Reduction half reaction: electrons go on the left (reactant)

o Half Reaction Method Steps:



1) Write separate equations for the oxidation and reduction half- reactions.

2) For each half reaction:

a. Balance all the elements except hydrogen and oxygen

b. Balance oxygen using

c. Balance hydrogen using H+

d. Balance charge using electrons

3) If necessary multiply one or both balanced half-reactions by an integer to equalize the

number of electrons transferred in the two half-reactions.

4) Add the half- reactions, and cancel identical species.

5) Check that the elements and charges are balanced.

XIV. Conversions and Stoichiometry

Look at the handout from teacher!!!

XV. Solutions and Molarity

Solutions:

o A solution is a mixture of two or more substances that are homogenous (the same

throughout)

o Solvent: the substance you have more of. The thing the substances are dissolved in. Usually

water.

o Solute: The substance you have less of. The substance that is dissolved.

o All solutions are made up of a solvent and a solute.

o Solvents and solutes can be liquids, solids, or gasses.

Ex) Salt water: Solvent- H2O (water) Solute- salt

Ex) Soda: Solvent- H2O (water) Solute: sugar, dye, CO2 (carbon dioxide)

Concentration: amount of solute dissolved in an amount of solvent

o Can be measured many different ways (ppm, Molarity) but the most common is molarity

(M)

Moles: of solute Liters: of solution

o Calculating Molarity:

Ex) 0.10 mol. of NaCl is dissolved in 1.0 L of H20. What is the molarity of the solution?

Ex) What is the concentration of a solution where 16.0g of sugar, , is dissolved in

250 mL of water?

TIP (for above problem): Always divide by volume.

Saturated Solution: Maximum amount of solute possible has dissolved in a solvent.

Unsaturated Solution: Less than the maximum has been dissolved.

Supersaturated Solution: More than twice the maximum has been dissolved.



A more concentrated solution means that there are more molecules per unit volume. Two

solutions with the same concentration, but with different volumes have the same number of

molecules per unit volume, but a different total number of molecules.

Preparing Solutions

o Two Methods for making a solution:

1) From a solid ( dissolve solid in H2O)

2) Diluting an already made solution

o Preparing from a solid (1)

Ex) you want to make a 0.15 M NaCl solution and you need 2.0L of that solution. What

is the mass of NaCl you will need?

** M*V(in liters)= Mol. g.

XVI. Gases

Kinetic Molecular Theory (KMT): How we model (predict) the behavior of gases in our

calculations.

A gas that follows the assumptions of KMT is called an ideal gas.

Many gases don’t follow KMT, those gases are called real gasses.

Assumptions of KMT:

1. Gases consist of a large number of molecules in constant, random motion.

2. The volume (size) of molecules is negligible (irrelevant) compared to the total volume of the

container.

3. The attractiveness and repulsive forces between molecules is negligible.

4. Collisions of gases are perfectly elastic (no sticking together, no energy lost)

5. The average Kinetic Energy (KE) of the molecules is proportional to the temperature of the

sample (definition of temperature)

At low temperature and high pressure the assumptions of KMT stop working (#2 and #3 are the

most important)

Relating Pressure (P) to KMT

o As long as the volume is held constant, pressure is dependent on only two things:

1. Temperature (avg. KE)

2. # of molecules

o Pressure is defined as the force exerted over and area ( P= F/a)

o What are the only two ways to increase the force?

1. Hit the sides more often (increasing the number of collisions/molecule)

2. Hit the sides harder (increasing KE/T)

o Pressure doesn’t depend on the mass of the molecule, if it did all the gas laws wouldn’t

work

3 Gas Laws

o Boyle’s Law: relates Pressure and volume. V

(volume is inversely related to pressure)



o Charles Law: relates volume and temperature. V T (volume is directly related to

temperature)

o Avogadro’s Law: relates volume and amount (moles). V n (volume is directly related to

amount)

Pressure Conversions

o 1atm= 760 mmHg = 760 torr = 101.352 kPa

o Pascals (Pa) is the SI unit for pressure

o Atmospheres (atm) is the most common SI Unit.

Temperature conversations: ALWAYS CONVERT TO KELVIN (K)!!!!!

o K= + 273.15

o Kelvin is always positive because absolute zero (0K) is the lowest possible temperature.

Standard temperature Pressure (STP)

o An arbitrarily chosen temperature and pressure used to compare gases.

o SI units: 273 K and 101.325 kPa

o Non-SI units: and 1 atm.

o BOLDED: what we use.

o If you have 1 mole of gas (any gas) if the gas is at STP it occupies a volume of 22.4 L

1 mol. = 22.4 L @ STP

Combined gas law

o Combines the three individual gas laws, and is used to determine what the new conditions

will be after a change.

o

o TIP: Cancel out the constants.

o TIP: Units don’t matter as long as they are the same on both sides of the equation.

(exception: temperature must by K (Kelvin))

PV=nRT

P- pressure (units- atm)

V- volume (units- L)

n- number of moles (units- moles)

R- gas constant (Units- L*atm/mol.*K)

T- temperature (units- K)

There are many R’s depending on the units (for this class: R= 0.0821 L*atm/mol.*K)

We use the ideal gas law to calculate all the variables of a gas.

Ex 1) What is the volume occupied by 0.118 moles of He at a pressure of 0.97 atm. And a

temperature of 305 K?

PV=nRT

TIP: solve for the desired variable before solving for the answer.

Ex 2) Would the volume be different (referring to the previous problem) if the gas were Ar

instead?

No, the identity and mass of a molecule doesn’t change the ideal gas law.



Ex 3) How many moles of air are in a 0.5L breath on top of Mt. Everest if the pressure is 0.33 atm

and the temperature is 254 K?

PV=nRT

Density= mass/volume. The density of a gas at STP can be defined as D= mass/volume.

For gases the units are g/L

Examples

Ex 1) what is the density of helium at STP?

Ex 2) what is the density of nitrogen gas at STP? (remember diatomic elements: H,N,F,O,I,Cl,Br)

The density of a gas is directly proportional to its molar mass. (i.e. the larger the mass the more

dense it is!)

Less dense gases float on top of more dense gases (air= 28.8 g/mol)

What if the gas isn’t at STP?

We can use the above equation to do the following

(1) Calculate the density of a gas under any conditions

(2) Calculate the molar mass of an unknown gas just by measuring its density.

Stoichiometry with gases involves converting volume of a gas to moles using:

1. PV=nRT (for any conditions)

2. 1 mol. = 22.4 L (if at STP)

If at STP you can convert from the volume of a to the volume of b just using the mole ratio

For examples see HW (the really long one )

XVII. Heating Curves

SEE ATTACHMENT 3 FOR DIAGRAM REFERED TO IN FOLLOWING BULLETS!!!

Definition of Temperature: measurement of the average kinetic energy of a substance.

o Think movement

o At points (a), (b), and (c) the thermal energy is being used to increase the temperature of

the substance (makes things move faster)

o At points (d) and (e) the thermal energy is being used to change the phase of the substance

and not to raise the temperature. The interactions of two separate molecules must be

broken for a phase change to occur.

o Where energy goes: on the diagonal line it goes to help the change in temperature. On the

flat line part it goes to help the change of state.

Conclusions from lab

o Heat transfer doesn’t always result in a change in temperature

o Heat of vaporization ( ): the energy required to vaporize a substance (liquid to gas)

o Heat of Fusion ( ): The energy requires to melt a substance ( solid to liquid)



o vs. : must completely separate molecules to boil, but just have “disorganize a

solid” to melt it.

XVIII. Thermodynamics Basics

Temperature: Heat is a transfer of thermal energy. Temperature is the measurement of average

kinetic energy.

First Law of Thermodynamics: The heat energy lost by one body is gained by another body. The

energy must be conserved.

Endothermic: Heat transfers from surroundings into the system. The system or reaction feels

cold. H is positive and q is positive.

Exothermic: Gives off energy, heat, and the heat transfers from the system to the surroundings.

The system or reaction feels hot and q is negative.

XIX. Calorimetry

o q= heat (J or cal.)

o m: mass (g)

o c: specific heat (J/g* c)

o delta T: change in temperature (units don’t matter)

XX. Kinetics

Definition: The study of the rates of reactions

o Rate: the change in a variable per unit time

Factors that affect reaction rates

o Collision theory: explains how reactions occur

1. Molecules must collide to react

2. Molecules need a minimum amount of energy to successfully react

Activation energy ( ): the energy required to transform reactions into

products (also called activation barrier)

3. Molecules must be aligned correctly for a reaction to occur

Factors that can control the rate of reaction

1. Catalyst: a substance that assists a reaction but isn’t consumed/changed by the reaction

A catalyst lowers the activation energy of a reaction by producing an alternate pathway

Connecting to collision theory: catalysts usually align molecules so that a successful

collision is more likely

2. Temperature: a higher temperature increases the rate of a reaction

Connections to collision theory:

a) Molecules collide more often because they are moving faster

b) Molecules have more energy so more collisions have the necessary activation

energy

3. Concentration: a higher concentration of molecules per unit volume increases the rate of

reaction.



Connection to collision theory: more molecules in the same volume increases the

probability of a collision

4. Particle size/surface area: more surface area increases the rate of a reaction (chunk of

marble vs. powder)

Chunk of marble: lower surface area, larger particle size

Powder: higher surface area, smaller particle size

Connection to collision theory: more molecules are exposed which means there is

more possibility for a collision/reaction.

o Reaction can be modeled using an energy diagram (see attachment 1)

o When (the change in enthalpy, heat energy) is positive the reaction is endothermic

(gets colder)

o When is negative the reaction is exothermic (gets hotter)

o DIAGRAM: SEE ATTACHMENT 4

o When is positive, reactants have less energy than the products (surroundings give

energy to reaction)

o When is negative, reactants have more energy than the products (surroundings absorb

energy from the reaction)

XXI. Nuclear Chemistry

See Reactivity Handout

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Study Guide Chemistry Comprehensive Exam 2014 I