Worksheet: Acids, Bases, and Salts Review



Acids, Bases, and Salts

Name______________

• An Arrhenius acid is defined as any compound that dissociates in aqueous solution to form ____________ ions.

HNO3( H+ + NO3

HCl (aq) ( ________________________

• An Arrhenius base is defined as any compound that dissociates in aqueous solution to form ____________ ions.

KOH (aq)( K+ (aq) + OH- (aq)

NaOH (aq)( ________________________

• Salts are compounds that dissociate in aqueous solution releasing neither ____________ ions nor ____________ ions.

KCl (aq) ( K+ (aq) + Cl- (aq)

NaCl (aq) ( ________________________

Using the Arrhenius definition, classify the following examples as acids, bases, or salts:

HBr ____________________ KCl __________________

Mg(OH)2 ________________ H3PO4 ________________

HCl ____________________ HClO _________________

KNO2 ___________________ Al(OH)3 ________________

HFO4 ___________________ KC2H3O2 _______________

Ba(OH)2 _________________ NaCl __________________

Acids and bases can also be identified using an operational definition. Operational definitions are simply a list of properties.

ACIDS:

♦ A ____________ taste is a characteristic property of all acids in aqueous solution.

♦ Acids react with some metals to produce ____________ gas.

♦ Because aqueous acid solutions conduct electricity, they are identified as ____________.

♦ Acids react with bases to produce a ____________ and water.

♦ Acids turn ____________ different colors.

BASES:

♦ Bases tend to taste ____________ and feel ____________.

♦ Like acids, aqueous basic solutions conduct ____________, and are identified as

____________.

♦ Bases react with ____________ to produce a salt and ____________.

♦ Bases turn ____________ different colors.

Naming Acids

• Binary acids consist of ____________ elements, the first being ____________.

• Ternary acids consist of ____________ elements. Do NOT use a prefix.

–ate becomes _______ and –ite becomes _______

Give the word equation for the neutralization reaction of an acid and a base.

Chemical Quantities Worksheet Name ______________________________

Period _____ Date ___________________

Recipes often specify the number of eggs needed. Although eggs are used individually when cooking small quantities, they are sold by the dozen, or by the gross (144, a dozen dozens), depending on the quantity wanted. If you were cooking for an army, you'd likely mix your eggs by the gross, rather than counting them individually. Instead of thinking in terms of 2 eggs per potato, you'd use 2 gross of eggs per gross of potatoes.

If you were buying rice for a casserole, would you go to a store and ask for 250,000 rice grains? Or would you ask for a pound of rice? Because rice is small, it's convenient to use other means of measuring than counting.

Chemical particles (atoms, molecules, etc.) are much, much, smaller than eggs or rice. It is therefore more convenient and useful to specify quantities in ways other than by counting individual atoms or molecules. The quantity called the mole is used to specify the number of particles, just like the dozen or gross is used for eggs and other items we encounter in our daily lives. Whereas a dozen is 12 of something, and a gross is 144 of something, a mole is 6.02 x 1023 of something (602,000,000,000,000,000,000,000). If a chemical reaction involves one atom of magnesium and two units of hydrochloric acid, one mole of magnesium will react with two moles of hydrochloric acid. We can't pick out individual chemical particles, but we can easily measure a mole of them.

How do we measure moles? We use a balance to determine the mass, and then convert mass to moles. Just as the mass of a gross of eggs differs from the mass of a gross of potatoes, the mass of a mole of aluminum atoms differs from the mass of a mole of calcium. The periodic table tells us the mass of one mole of each element. For instance, a mole of aluminum has a mass of 26.98 grams. A mole of calcium has a mass of 40.08 grams.

Use the periodic table to determine the mass of the following quantities of chemical substances. Remember, the periodic table tells you the MOLAR MASS, the mass of one mole of each element. The mass of two moles would be twice that of one mole. For compounds, add the molar masses in the ratios indicated in the chemical formulas for the molar mass of the compound.

1 mole of lithium _______________

1 mole of magnesium _______________

1 mole of carbon _______________

1 mole of oxygen _______________

2 moles of oxygen _______________

1 mole of hydrogen _______________

2 moles of hydrogen _______________

1 mole of H2O _______________

1 mole of CO2 _______________

2 moles of CO2 _______________

Use the periodic table to determine the number of moles in the following masses. Remember, the periodic table tells you the MOLAR MASS, the mass of one mole of each element. For compounds, add the molar masses in the ratios indicated in the chemical formulas for the mass of one MOLE of the COMPOUND. Divide the mass indicated by the molar mass to calculate the number of moles.

35.45 grams of chlorine _______________

63.55 grams of copper _______________

18.024 g of beryllium _______________

36.033 grams of carbon _______________

100 grams of hydrogen _______________

100 grams of iron _______________

100 grams of lead _______________

18.0 grams of H2O _______________

36.0 grams of H2O _______________

100 grams of H2O _______________

1. Balance the following equation: ____ K2PtCl4 + ___ NH3 → ___ Pt(NH3)2Cl2 +___ KCl

Determine the grams of KCl produced (theoretical yield) if you start with 34.5 grams of NH3.

2. Balance the following equation: H3PO4 + 3 KOH → K3PO4 + 3 H2O

If 49.0 g of H3PO4 reacts, how many grams of K3PO4 should be produced?

3. Balance the following equation: Al2(SO3)3 + 6 NaOH → 3 Na2SO3 + 2 Al(OH)3

If you start with 389.4 g of Al2(SO3)3 and how many grams of Na2SO3 would be produced?

4. Balance the following equation: Al(OH)3 (s) + 3 HCl (aq) → AlCl3 (aq) + 3 H2O (l)

If you start with 50.3 g of Al(OH)3 how many grams of AlCl3 would you expect to produce? Assume you did the experiment and you only got 39.5 g of AlCl3, what is the percent yield (hint: think about how you determine your % for a test – what you got divided by what you should have gotten)?

5. Balance the following equation: K2CO3 + 2 HCl → H2O + CO2 + 2 KCl

Determine the theoretical yield (grams) of H2O if you start with 34.5 g of K2CO3. If only 3.4 g of H2O is produced, what is the percent yield?

6. Balance the following equation: H2SO4 + Ba(OH)2 → BaSO4 + 2 H2O

a) If 98.0 g of H2SO4 reacts, how many grams of BaSO4 should be produced? b) What is the percent yield if you only produce 213.7 g of BaSO4? (use the answer from a to determine the answer for b)

Gas Laws

SHOW ALL WORK FOR ALL PROBLEMS

I. 1.0 atm = 101.3 kPa = 760 mmHg And 0(C = 273 K

Change the following units: 359 kPa = _________ atm 10(C = ________ K

6.2 atm = ________ kPa 10K = _______ (C

For the rest of the problems: First identify each number with P, V, or T. Second state whose law you are using, Third – show the equation, Fourth solve the problem, and Fifth - circle your final answer - and make sure you don't forget your units!!!

1. The gas in a sealed can is at a pressure of 3.00 atm at 25(C. A warning on the can tells the user not to store the can in a place where the temperature will exceed 52(C. What would the gas pressure in the can be at 52(C?

2. A sample of hydrogen exerts a pressure of 0.329 atm at 47(C. The gas is heated 77(C at constant volume. What will its new pressure be?

3. 3. A sample of neon gas occupies a volume of 752 mL at 25(C. What volume will the gas occupy at standard temperature if the pressure remains constant?

4. A sample of oxygen gas has a volume of 150 mL when its pressure is 440 mmHg. If the pressure is increased to standard pressure and the temperature remains constant, what will the new gas volume be?

5. Ral3ph had a helium balloon with a volume of 4.88 liters at 150 kPa of pressure. If the volume is changed to 3.15 liters, what would be the new pressure in atm?

6. 5.36 liters of nitrogen gas are at -25(C and 733 mm Hg. What would be the volume at 128(C and 1.5atm?

7. At constant temperature, 2 L of a gas at 4 atm of pressure is expanded to 6 L. What is the new pressure? (Do this one conceptually and not algebraically.)

Thermochemistry

Specific Heat Worksheet

C = q/m∆T, where q = heat energy, m = mass, and T = temperature Remember, ∆T =

(Tfinal – Tinitial). Show all work and proper units.

1. A 15.75-g piece of iron absorbs 1086.75 joules of heat energy, and its temperature

changes from 25°C to 175°C. Calculate the specific heat capacity of iron.

2. How many joules of heat are needed to raise the temperature of 10.0 g of aluminum

from 22°C to 55°C, if the specific heat of aluminum is 0.90 J/g°C?

3. To what temperature will a 50.0 g piece of glass raise if it absorbs 5275 joules of heat

and its specific heat capacity is 0.50 J/g°C? The initial temperature of the glass is

20.0°C.

4. Calculate the heat capacity of a piece of wood if 1500.0 g of the wood absorbs

6.75×104 joules of heat, and its temperature changes from 32°C to 57°C.

5. 100.0 mL of 4.0°C water is heated until its temperature is 37°C. If the specific heat of

water is 4.18 J/g°C, calculate the amount of heat energy needed to cause this rise in

temperature.

What is enthalpy?

What is enthalpy?

What is Hess’s Law?

Find the ΔH for the reaction below, given the following reactions and subsequent ΔH values: 

PCl5(g)  →  PCl3(g)  +  Cl2(g)

P4(s)  +  6Cl2(g)  →  4PCl3(g)            ΔH = -2439 kJ 

4PCl5(g)  →  P4(s)  +  10Cl 2(g)         ΔH = 3438 kJ 

(2)  Find the ΔH for the reaction below, given the following reactions and subsequent ΔH values: 

2CO2(g)  +  H2O(g)  →  C 2H2(g) +  5/2O2(g)

C2H2(g) + 2H2(g)  →  C2H6(g)                              ΔH  =-94.5 kJ 

H2O(g)  →  H2(g) + 1/2O2 (g)                               ΔH  =71.2 kJ 

C2H6(g) +  7/2O2(g)  →  2CO2(g)  +  3H2O(g)     ΔH  =-283 kJ 

3)  Find the ΔH for the reaction below, given the following reactions and subsequent ΔH values: 

N2H4(l)  +  H2(g)  →  2NH3(g)

N2H4(l)  +  CH4O(l)  →  CH2O(g)  +  N2(g)  +  3H2 (g)         ΔH = -37 kJ 

N2(g)  +  3H2(g)  →  2NH 3(g)                                                ΔH = -46 kJ 

CH4O(l)  →  CH2O(g) +  H 2(g)                                              ΔH = -65 kJ 

(4)  Find the ΔH for the reaction below, given the following reactions and subsequent ΔH values: 

H2SO4(l)  →  SO3(g)  +  H2O(g)

H2S(g)  +  2O2(g)  →  H2SO4(l)                                  ΔH = -235.5 kJ 

H2S(g)  +  2O2(g)  →  SO 3(g)  +  H2O(l)                    ΔH = -207 kJ 

H2O(l)  →  H2O(g)                                                      ΔH = 44 kJ 

(5)  Find the ΔH for the reaction below, given the following reactions and subsequent ΔH values: 

2C2H4O(l) + 2H2O(l)  →  2C2H6O(l) +  O2(g)

C2H6O(l)  +  3O2(g)  →  2CO2(g)  +  3H2O(l)            ΔH = -685.5 kJ 

C2H4O(l)  +  5/2O2(g)  →  2CO2(g)  +  2H2O(l)         ΔH = -583.5 kJ 

Nuclear Chemistry

What is nuclear fission?

What is nuclear Fusion?

Where is Nuclear Fusion used?

What are Isotopes?

What is a chain reaction?

Write the equations for:

1) the alpha decay of radon198

2) The beta decay of uranium 237

3) Positron emission from sodium 22

4) Write the symbols for an alpha particle, beta particle, gamma ray, and

positron.

5) If the half-life for the radioactive decay of Mendelevium 101 is 5 minutes and I

start with a 130 gram sample, how much will be left over after 60 minutes?

Chemistry Semester Two Key Term List

These are the terms you are expected to know the meaning of when read, by the time you have completed Spring Chemistry.

Molarity solute

Solvent

boiling point elevation Brownian motion

colligative property concentration

dilution freezing point depression

heat of solution hydrogen bonding

immiscible insoluble

miscible osmotic pressure

osmosis saturated

solubility solution

supersaturated unsaturated

amphoteric neutralization

Acid Anhydride

Arrhenius acid/base Base

Bronsted-Lowry model buffer

conjugate acid/base equivalence point

Hydronium ion Indicator

Ionization pH, pOH

salt standard solution

strong acid/base titration

weak acid/base

Ka and Kb Lewis acid/base

Kelvin Standard pressure

Standard temperature

atmospheric pressure Avogadro’s principle

Barometer Boyle’s Law

Charles’s Law Combined Gas Law

Gay-Lussac’s Law Ideal Gas Law

Kinetic molecular theory Pascal

enthalpy specific heat

activated complex activation energy

calorie catalyst

collision theory energy

entropy heat

heat of formation heat of reaction

inhibitor intermediate

Joule law of conservation of energy

reaction rate spontaneous process

thermochemical equation chemical potential energy

Gibbs Free Energy

spontaneous reaction state functions

fission fusion

alpha particle atom

atomic mass atomic mass unit

atomic number beta particle

chain reaction electron

gamma radiation half-life

isotope mass number

neutron nuclear reaction

nucleus positron

proton radiation

radioactive decay

Functional group Organic molecule

Alcohols Aldehydes

Alkane Alkene

Alkyl Halides Alkyne

Amines Aromatic

Carboxylic acids Cycloalkanes

Esters Ethers

Geometric isomers Ketones

Monomer Polymer

Saturated hydrocarbons Structural isomers

Unsaturated hydrocarbon

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