Chemistry Final Review

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Chemistry Final Review

2^nd Semester, 2015-2016

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Unit 1: Basics

• How many sig figs are in the following numbers?

• 2300

• 2

• 314

• 3

• 2.06

• 3

• 0.0025

• 2

• 9.001

• 4

• 6.02x10²³

• 3

• 587.0

• 4

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Unit 1: Basics

Discuss the following in terms of accuracy

and precision

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Unit 1: Basics

• List the three mole conversions you learned at the beginning of the year.

• 1 mol = 6.02x10²³ particles (atoms, molecules, ions, etc)

• 1 mol = molar mass in grams from the periodic table

• 1 mol = 22.4 L of GAS at STP see how that makes more sense now ☺?

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Unit 1: Basics

• Write the following in standard or scientific notation

• 2.71x10⁴

• 27100

• 6.4x10^-3

• 0.0064

• 2687

• 2.687x10³

• 0.012654

• 1.2654x10^-2

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Unit 1: Basics

• What is the equation for density?

• Density = mass volume

• What are the common units for density?

• g/mL or g/cm³

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Unit 1: Basics

• Give an example of an element and an example of a compound.

• Element: He, C, Mn

• Compound: CO_2,NaCl, etc (anything with more than one element)

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Unit 1: Basics

• Write the equation for percent yield.

• % yield = actual x 100 expected

• Write the equation for percent error.

• % error = (actual-expected) x 100 expected

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Unit 2: Energy Transfer

• Draw a heating curve for water.

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Unit 2: Energy Transfer

• Label the Q-equations for each section.

Q=mcΔT

Q=mHv

Q=mHf

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Unit 2: Energy Transfer

• Find the value for Hf, c, and Hv in your data book.

• Hf = 79.72 cal/g

• C = 1.0 cal/g°C

• Hv = 539.4 cal/g

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Unit 3: Atoms & Periodic Table

• Complete the following table:

Element/

ion

Atomic number

Atomic mass

Protons Neutrons electrons

Fe Cl^- K⁺

26 55.85 26 30 26

17 35.45 17 18 18

19 39.10 19 20 18

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Unit 3: Atoms & Periodic Table

• Positive ions form when:

• Atoms lose electrons (usually metals, on left of table)

• Negative ions form when:

• Atoms gain electrons (usually nonmetals, on right of table)

• Why do atoms form ions?

• To become more stable, get the configuration of a noble gas.

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Unit 3: Atoms & Periodic Table

• The halogens make a charge of ____ when they become ions.

• -1

• The alkali metals make a charge of ___ when they become ions.

• +1

• The alkaline earth metals make a charge of ___

when they become ions.

• +2

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Unit 3: Atoms & Periodic Table

• What are the three types of nuclear decay?

• Alpha, Beta, Gamma

• What type of particle does each emit?

• Alpha = helium nucleus (2 protons, mass of 4)

• Beta = electron (no mass, -1 proton)

• Gamma = high energy (no mass, no proton change, but product is more stable)

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Unit 3: Atoms & Periodic Table

• Complete the following:

• Type of Decay

• ________ ^99m₄₃Tc ⁹⁹₄₃Tc + ______

• ________ ²⁴⁷₉₅Am ⁰_-1e + _____

• ________ ¹⁷⁵₉₃Np ⁴₂He + ____

0 0γ

247

96Cm

171

91Pa Gamma

Beta Alpha

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Unit 4: Atoms & Periodic Table

• Draw a Bohr model for Beryllium.

• Draw a Bohr model for Silicon.

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Unit 4: Compounds & Bonding

• Which types of elements participate in ionic bonding?

• Metals (+) and non-metals (-) (also polyatomic ions)

• Which types of elements participate in covalent bonding?

• Non-metals and non-metals (they share

electrons instead of charges sticking together)

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Unit 4: Compounds & Bonding

Name the following compounds:

• MgO

• Magnesium oxide

• AlF₃

• Aluminum fluoride

• NiSO₄

• Nickel (II) sulfate

• FeCl₂

• Iron (II) chloride

• N₂O₅

• Dinitrogen pentoxide

• SF₄

• Sulfur tetrafluoride

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Unit 4: Compounds & Bonding

• Describe a polar covalent bond.

• Covalent bond (non-metal and non-metal) in which electrons are shared UNEVENLY (one atom is more electronegative than the other).

• Draw a water molecule and show its polarity.

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Unit 4: Compounds & Bonding

• List and describe the 4 types of Intermolecular forces (IMF).

• Dispersion forces (weakest, between nonpolar molecules)

• Dipole-dipole interaction (stronger, between polar molecules)

• Hydrogen-bonding (STRONG, between molecules with N, O, or F bonded to H, explains water’s high boiling point)

• Ion-molecule interaction (strongest, how water dissolves salt – pulls apart the ions)

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Unit 5: Chemical Reactions

• Write the generic equation for all 6 types of chemical reactions

• Synthesis: A + B AB

• Decomposition: AB A + B

• Single Replacement: A + CB AB + C (remember, there are 3 sub-types of Single Replacement)

• Double Replacement: AB + CD AD + CB

• Combustion: Fuel + __O₂ __CO₂ + __H₂O

• Acid/Base: HD + COH CD + H₂O

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Unit 5: Chemical Reactions

• What table do you use to decide if a single

replacement reaction happens? What do you look for?

• Table N

• Strong metals are on the bottom…if the LONE metal is closer to the bottom it will replace the metal in a compound.

• Strong halogens are at the top…if the LONE halogen is closer to the top, it will replace the halogen in a compound.

• Hydrogen follows the metal rule.

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Unit 5: Chemical Reactions

• List the 7 diatomic elements.

• H, O, N, Cl, Br, I, F

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Unit 5: Chemical Reactions

• What table do you use to decide if a double replacement reaction happens? What do you look for?

• Table E (solubility)

• Both reactants must be aqueous (soluble), AT LEAST one of the products must be NOT

AQUEOUS (solid, liquid, or gas)

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Unit 5: Chemical Reactions

Single and Double Replacement – predict the products and write the balance equation

1. silver nitrate + nickel

2AgNO₃ (aq) + Ni (s) Ni(NO₃)₂ (aq) + 2Ag (s)

2. lead + zinc acetate Pb + Zn(C₂H₃O₂)₂ N.R.

3. NaOH + CaBr₂

2 NaOH (aq) + CaBr₂ (aq) 2 NaBr (aq) + Ca(OH)₂ (s) 4. Pb(NO₃)₂ + HCl

Pb(NO₃)₂ (aq) + 2 HCl (aq) 2 HNO₃ (aq) + PbCl₂ (s)

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Unit 5: Chemical Reactions

• 1. Consider this molecular (normal) equation:

H₃PO₄ (aq) + 3KOH (aq) K₃PO₄ (aq) + 3H₂O (l)

2. Write the complete ionic equation:

3H⁺(aq) + PO₄^3-(aq) + 3K⁺(aq) + 3OH^-(aq) 3K⁺(aq) + PO₄^3-(aq) + 3H₂O(l) 3. Write the net ionic equation:

3H⁺(aq) + 3OH^-(aq) 3H₂O(l) 4. What are the spectator ions?

PO₄^3-(aq), 3K⁺(aq)

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Unit 6: Stoichiometry

49.

In ammonia production, nitrogen and hydrogen are synthesized into ammonia (NH

₃

). What mass of

ammonia will be produced from 1.5 kg of nitrogen

assuming that hydrogen is in

excess? (Hint: first write the

complete, balanced equation)

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49A.

N

₂

(g) + 3H

₂

(g) 2NH

₃

(g)

1.8 kg NH

₃

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Unit 6: Stoichiometry

50. A solution made from 5.0 g of

copper (II) sulfate is mixed with a solution containing excess calcium nitrate. A precipitate of calcium sulfate is formed.

a) Write the complete, balanced equation for the reaction.

b) Write the complete ionic equation for the reaction.

c) Write the net ionic equation for the

reaction.

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Unit 6: Stoichiometry

50. A solution made from 5.0 g of

copper (II) sulfate is mixed with a solution containing excess calcium nitrate. A precipitate of calcium sulfate is formed.

Continued:

d) How much (mass) calcium sulfate is expected?

e) If the amount of CaSO

₄

measured in the experiment was 3.99 g, what is the percent error?

f) What is the percent yield?

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50A.

a) CuSO_{4 (aq)} + Ca(NO₃)_{2 (aq)} CaSO_{4 (s)} + Cu(NO₃)_{2 (aq)} b)Cu²⁺+SO₄^2-+ Ca²⁺+ 2NO₃^- CaSO_{4 (s)} +Cu²⁺+ 2NO₃^-

c)SO

₄^2-

+ Ca

²⁺

CaSO

₄

(s) d)4.3 g CaSO

₄

e)-7.2%

f)93%

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Unit 7: Solutions

Use Table E to determine the solubility of each

substance:

ammonium chloride barium carbonate

silver iodide

mercury (II) bromide

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ammonium chloride

^soluble

barium carbonate

nearly insoluble

silver iodide

nearly insoluble mercury (II) bromide

slightly soluble

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Unit 7: Solutions Use Table D:

1. How many grams of sodium nitrate will dissolve in 100 g of water at 25 C?

2. How many grams of ammonia (NH₃) will dissolve in 100 g of water at 100C?

3. If 140 g of KI is dissolved in 100 g of water at 30 C, is the solution saturated,

Chemistry Final Review