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AcademicAcademic ChemistryChemistryUNIT 7UNIT 7

PERIODIC TRENDSPERIODIC TRENDS& CHEMICAL& CHEMICALEQUATIONSEQUATIONS

Name:

Class Period: Test Date:

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Chemistry Calendar

Monday Tuesday Wednesday Thursday FridayJanuary 2nd

Teacher Prep Day

3rd

Teacher Prep Day

4th

Notes #1: History and Review of the Periodic Table

5th

Notes #2: Definitions and Trends

6th

Graphing Activity

9th

QUIZNotes #3: Octet Rule and Valence e-‘s

10th

Notes #4: Bonding Review and Intro to RxnsConservation of Matter DEMO

11th

Notes #5: Rxns (synthesis and decomp)Decomposition DEMO

12th

Notes #5: Rxns (Single and Double Replacement)CuCl2 + Al DEMO

13th

Notes #5: Rxns (RedOx and Combustion)Electrolysis Lab

16th

MLK DayStudent Holiday

17th

Reactions PracticeSingle Replacement Rxn Lab

18th

Benchmark/Test ReviewDouble Replacement lab?

19th

Benchmark/Test ReviewDouble Replacement lab?

20th

TEST

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NOTES #1: Textbook Activity

Use the textbook to fill in the notes about the following main ideas.

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MAIN IDEA NOTESDmitri Mendeleev

. .

Henry Mosley

. .

Periodic Law

. .

Period

. .

Group

. .

Metals

. .

Alkali Metals

. .

Alkaline Earth Metals

Mendeleev organized elements according their chemical and physical properties.

Arrangement allowed him to predict the existence and properties of three missing elements (Ga, Sc, Ge)

Periodic: repeating pattern Mendeleev’s periodic table is arranged in order of increasing atomic mass.

. . Moseley rearranged Mendeleev’s periodic table.

His periodic table is arranged in order of increasing atomic number.

This is the way our current periodic table is still organized.. . A periodic repetition of chemical and physical properties of the elements

when they are placed in order of increasing atomic number.

. . Row of the periodic table. Periods have properties that change as you move across the row.. . column of the periodic table Groups have similar chemical and physical properties.. . Most of the elements on the periodic table are metals. Located to the left of the zigzag line. All are solids at room temperature except Mercury (Hg). Gallium (Ga) is the

only metal that is a liquid at human body temperature.

. . Traits: Soft, silvery metals that have a low density. They are the most

reactive metals and are not found by themselves in nature (they are always in compounds). Explosive when exposed to air and water.

Uses: Used in soap, lights, and are present in the body (electrolytes). Examples: Li, Na, K, Rb, Cs, Fr

. . Traits: Harder, denser and stronger than the Alkali metals. Silvery in color.

They are the second most reactive metals and are not found by themselves in nature (they are always in compounds).

Uses: Gemstones, limestone, marble, alloys for aircraft, truck bodies, fireworks and flares. Calcium needed for strong teeth and bones.

Examples: Be, Mg, Ca, Sr, Ba, Ra

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MAIN IDEA NOTESTransition Metals

. .

Inner Transition Metals

. .

Metalloids

. .

Noble Gases

. .

Nonmetals

. .

Halogens

Traits: Hard, silvery, solid metals that have a high density. High luster (shiny), malleable, ductile and good conductors of heat and electricity. They are less reactive than alkali and alkaline earth metals and many are found by themselves in nature.

Uses: Paints, steel, coins, electrical wiring, plumbing, jewelry, construction. Examples: Fe, W, Cu, Ag, Au, Pt, Ni

. . Traits: Shiny, silvery, reactive metals that have a high melting point.

Lanthanides are usually found in compounds. Actinides are all radioactive (unstable). Many actinides are artificially produced. The only actinides found in nature are Th, Pa, U and Np.

Uses: Lanthanides are used in tv tubes. Examples: Ce, Pm, Eu, Am, U, Es

. . Traits: Properties of both metals and nonmetals. Semiconductors. Good

electrical conductors at high temperature, good insulators at low temperature.

Uses: computer chips and glass Examples: B, Si, Ge, As, Sb, Te, Po, At

. . Traits: No chemical reactivity (inert). All are gases. Uses: Hot-air balloons Examples: He, Ne, Ar, Kr, Xe, Rn

. . Traits: Poor electrical conductors but are good insulators. They do not have

luster and are brittle in the solid form. Non-ductile. Exist in various phases (states) of matter: solids, liquids, or gases.

. . Fluorine and chlorine are yellow-green gases. Bromine is a dark-red liquid.

Iodine is purple-black crystalline solid Most reactive group of nonmetals. Uses: gemstones, bones, teeth, DNA, medicines, organic compounds, Xerox

copying, bleach, antiseptic

MAIN IDEA NOTES

NOTES #2: Definitions & Periodic Trends

WARM-UP: 1. As you move from left to right across the periodic table, the number of valence electrons increases

.

Atomic mass definition:

the weighted average of the masses (protons + neutrons) of the isotopes of an element.

TREND:

Atomic mass increases left to right

.

. A tomic mass increases

.

Atomic radius definition:

one half the distance between the nuclei in a molecule consisting of identical atoms. TREND:

Atomic radii decreases from left to right

.

. Atomic radii increases

.

Shielding Effect:

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REASON:

Adding more protons and neutrons!

REASON:

More protons to the right, increases nuclear charge on electrons pulling them closer to the nucleus. Therefore atomic radii decreases. Principal energy

Electrons in outer shells are repelled by electrons occupying the inner shells. They are not as strongly attracted to the positive nucleus as a result.

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Ionization Energy definition:

the energy required to remove an electron from a gaseous atom. First ionization energy- energy required to remove the outermost electron.TREND:

Ionization energy increases left to right

.

. Ionization energy decreases

.

Electronegativity definition:

Tendency for the atoms of an element to attract electrons when they are chemically combined with atoms of another element TREND:

Electronegativity increases left to right

.

. Electronegativity decreases

.

Reactivity definition:

Reactivity refers to how likely or vigorously an atom is to react with other substances. This is usually determined by ionization energy and electronegativity because it is the

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REASON:Left to right – Nuclear charge increases & shielding effect is constant resulting in greater attraction of the nucleus for the electron.Moving down –

REASON:

Left to right – the octet rule. More likely to attract electrons when combined if closer to 8. Moving down – valence electrons

transfer/interaction of electrons that is the basis of chemical reactions. Trend for metals: Trend for nonmetals:

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Reactivity decreases Re

activ

ity

incr

ease

s

Reac

tivity

de

crea

ses

Reactivity increases

ASSESSMENT: Periodic Trends

Matching (#1-4):

D 1. electronegativity

B 2. shielding effect

C 3. ionization energy

A 4. atomic radius

Which has the larger radius?

I 5. I or Xe

Si 6. Si or P

Br 7. F or Br

P 8. P or S

Tl 9. Al or Tl

Ar 10. Ne or Ar

Which is MORE reactive?

Fr 17. Fr or K

Ra 18. Be or Ra

F 19. I or F

O 20. O or Se

N/A 21. He or Ne

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Which has a LOWER ionization energy?

As 11. As or Se

Al 12. B or Al

S 13. S or Cl

Be 14. Be or B

At 15. O or At

K 16. K or Kr

A. The distance from the nucleus to the outermost electron of an atom.B. The reduction of the attractive force between a positively charged

nucleus and its outermost electrons due to the cancellation of some of the positive charge by the negative charge of the inner electrons.

C. The amount of energy needed to remove an outer electron from a specific atom or ion in its ground state (in the gas phase).

D. The tendency of atom to attract electrons to itself when bonded to another atom.

Which has the HIGHER electronegativity?

Ge 22. Ga or Ge

Br 23. Se or Br

Se 24. As or Se

N 25. N or P

O 26. O or I

K 27. K or Cs

ACTIVITY: GRAPHING PERIODIC TRENDS

PRE-LAB DISCUSSION:The Periodic Table is arranged according to the Periodic Law. The Periodic Law states that when elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern. Students can discover these patterns by examining the changes in properties of elements on the Periodic Table. The properties that will be examined in this lesson are: atomic radius AND first ionization energy.

PURPOSE: To understand trends of the periodic table and practice methods of graphing.

PROCEDURE: Graph the following information according to the steps described.

Symbol Atomic Number Atomic Radius (Picometers)

First Ionization(Energy-Joules)

Period 2Li 3 1.23 124 Be 4 0.89 215 B 5 0.80 191 C 6 0.77 260 N 7 0.70 335O 8 0.66 314 F 9 0.64 402

Ne 10 0.67 497 Period 3

Na 11 1.57 119 Mg 12 1.36 176 Al 13 1.25 138 Si 14 1.17 188 P 15 1.10 242 S 16 1.04 239 Cl 17 0.99 299 Ar 18 0.98 363

Period 4K 19 2.03 100

Ca 20 1.74 141 Rb 37 2.16 96 Sr 38 1.91 131 Cs 55 2.35 90 Ba 56 1.98 120

CALCULATIONS AND GRAPHS:DIRECTIONS: You will construct FOUR graphs, please read through the directions for each before beginning.

Graph 1For elements 3-20 make a graph of atomic radius as a function of atomic number. Plot atomic number on the X axis and atomic radius on the Y axis. After creating the graph, use a colored pen or pencil to draw a vertical line that represents that beginning of each period (horizontal row on the periodic table).

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Graph 2For elements in Group 1 (Alkali metals), make a graph of atomic radius as a function of atomic number. Make a second line on this same graph that will represent Group 2 (Alkaline Earth Metals). Use a periodic table to determine which elements are members of Group 1 and which elements are members of Group 2.

Graph 3For elements 3-20, make a graph of the energy required to remove the easiest electron (first ionization energy) as a function of atomic number. Plot atomic number on the X axis and energy required on the Y axis.After creating the graph, use a colored pen or pencil to draw a vertical line that represents that beginning of each period (horizontal row on the periodic table).

Graph 4For elements of Group 1 (Alkali metals), make a graph of the energy required to remove the easiest electron (first ionization energy) as a function of atomic number. On the same graph make a second line to represent Group 2 (Alkaline Earth Metals). Use a periodic table to determine which elements are members ofGroup 1 and which elements are members of Group 2.

QUESTIONS FOR DISCUSSION:1. What happens to the atomic radius as the atomic number increases across a period? Down a group?

2. What happens to the energy needed to remove an electron as the atomic number increases across a period? Down a group?

3. Why does atomic radius change as it does?

4. Why does the energy required to remove an electron change as it does?

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Graph 1

GRAPH 2

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Graph 3

Graph 4

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NOTES #3: The Octet Rule and Valence Electrons

WARM-UP: 1. How many valence electrons are in the outer shell of each of the following:

a. Cl - 7 c. Al - 3b. P – 5 d. Mg - 2

The Octet Rule:1. All atoms want 8 valence electrons (except for H and He – they want 2 valence electrons.) Atoms will gain, lose

or share electrons in order to get to 8 valence electrons. o The noble gases in group 8A (or 18) already have 8 electrons in the outer shell and are the most stable.

2. Electrons have a NEGATIVE charge.o If an atom loses electrons, it becomes more positive. These atoms are called Cations.o If an atom gains electrons it becomes more negative. These atoms are called Anions.

3. Valence electrons can be found by using the periodic table

1. How many valence electrons does Ca have? 2 What ion does it form? Ca 2+

2. How many valence electrons does Sr have? 2 What ion does it form? Sr 2+

3. How many valence electrons does Li have? 1 What ion does it form? Li +

4. How many valence electrons does Si have? 4 What ion does it form? Si 4-

5. How many valence electrons does Cl have? 7 What ion does it form? Cl -

6. How many valence electrons does H have? 1 What ion does it form? H +

***Do not forget that if it ends in “ate” or “ite” it is a polyatomic ion and IS NOT found on the periodic table***

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Bohr Models:

Electrons orbit in orbitals or more precisely – clouds. But for simplicity purposes, scientists draw electrons in shells.

• 1st shell can fit up to 2 electrons. Think about the electron configurations. This first shell is 1s.

• 2nd and 3rd shells can each fit up to 8 electrons. These are the 2s and 2p shells or the 3s and 3p shells, respectively.

The diagrams above are Bohr models. The show all of an atom’s electrons, but can take a while to draw – especially for larger atoms. Chemists prefer to draw Lewis Dot Structures, which only show valence electrons and the outer shell and are therefore quicker to draw and easier to read.

Lewis Dot Structures Review: EXAMPLES:

Practice: Draw the Lewis Dot Structures for the following atoms

1. He 3. Mg

2. Cl 4. N

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1st Period/Row

2nd Period/Row

3rd Period/Row

Directions:1. Start by writing the atomic symbol.2. Determine the number of VALENCE electrons.3. Draw dots for each of the valence electrons on each

of the four sides of the symbol.a. Put one dot on each side before you start to pair

the electrons

Notes #4: Bonding Review and Intro to Reactions

In a chemical reaction one or more bonds are formed, broken and reformed to create new substances. The original substances are referred to as the reactants and the resulting substances are the products.

Reactants Products

In the following examples, see if you can identify the reactants and the products as we review the types of bonds that can be formed.

Ionic bond: The donating/receiving of electrons to form a bond. This bond is between metals and nonmetals or metals and polyatomic ions.

Covalent bond: The sharing of electrons to form a bond. This bond forms between nonmetals and nonmetals.

Practice: Identify the following types of bonds (Ionic or Covalent)

1. CCl4 Covalent 7. MgO Ionic

2. MgCl Ionic 8. CH4 covalent

3. Na2NO3 Ionic 9. CaS ionic

4. C2H6O2 Covalent 10. HCl covalent

5. H2O Covalent 11. LiBr ionic

6. NH4Cl Covalent 12. K2SO3 ionic

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Na + Cl → Na Cl

Sodium DONATES one electron to chlorine leaving both atoms with full outer shells thus the octet rule is fulfilled.

Cl + Cl → Cl ClOne Chlorine atom SHARES an electron with the second Chlorine atom thus fulfilling the octet rule.

Chemical Reactions

Symbols: Explanation+ with, and; separates reactants/products yields, gives, produces, reacts to form(s) solid state (l) liquid state(g) gas state(aq) aqueous state; dissolved in water heat used in reaction Pt catalyst in reaction (ex. Platinum)

Reversible reaction

Skeletal equations show no relative amounts of reactants or products.

1. Mg (s) + H2SO4 (l) MgSO4 (aq) + H2

Word equations explain reactions using names.

2. Solid tetraphosphorus decoxide reacts with water to produce phosphoric acid.

Balancing Chemical Reactions

1. Count the number of each element for both sides of arrow.

2. Add (coefficient) whole number at beginning of each compound to balance elements.

3. Each element must have same number on both sides of arrow. All in smallest ratio.

4. If number is a subscript, multiply by coefficient to tell how many of each.

5. There are 7 diatomic molecules which when not combined with another element, written as below. H2 N2 O2 F2 Cl2 Br2 I2

6. When writing formulas from names, apply rules such as crisscross.

Guided Practice: Balance these equations:

1. F2 + KCl KF + Cl2

2. Al + CuSO4 Al2(SO4)3 + Cu

3. Aluminum + Oxygen Aluminum oxide

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1 2 2 1 **How many of each element do we have? Are both sides equal?

2 3 1 3

**Write out the equation using nomenclature and criss-cross and then balance.

4 Al + 3 O2 -> 2 Al2O3

Quick Nomenclature Review: Ionic Compounds: Name the metal unchanged, Name the nonmetal with “ide” on the end. Covalent Compounds: Use the prefixes listed below to name the compounds

1* 2 3 4 5 6 7 8 9 10Mono- Di- Tri- Tetra- Penta- Hexa- Hepta- Octa- Nona- Deca-

Compounds with Polyatomic Ions: Write both the metal and the polyatomic ions AS IS . Acids: Use the chart below

ION TYPE ION ENDING ACID NAME BEGINNING ACID ENDING

Polyatomic-ite NO hydro- beginning -ous-ate NO hydro- beginning -ic

Monatomic -ide hydro- beginning -ic

Independent Practice:

1. 2 C2H6 + 7 O2 4 CO2 + 6 H2O

2. Magnesium Nitrate + Calcium Hydroxide Calcium Nitrate + Magnesium Hydroxide **use polyatomic Ion list.

Mg (NO3)2 + Ca(OH)2 → Mg(OH)2 + Ca(NO3)2 Balanced

3. _____ H2 + _____ Cl2 2 HCl

4. Magnesium Chloride + Sodium Iodide Magnesium Iodide + Sodium Chloride

MgCl2 + 2 NaI → MgI2 + 2 NaCl

5. 2 HgO 2 Hg + _____ O2

6. Aluminum Carbide + Sodium Oxide Aluminum Oxide + Sodium Carbide

Al4C3 + 6 Na2O → 2 Al2O3 + 3 Na4C

7. _____ Ca + 2 H2O _____ Ca(OH)2 + _____ H2

8. _____ CH4 + 2 O2 _____ CO2 + 2 H2O

9. _____ Na2O2 + _____ H2SO4 _____ Na2SO4 + _____ H2O2 balanced

10. Copper (II) Carbonate Copper (II) Oxide + Carbon Dioxide

CuCO3 → CuO + CO2 Balanced

11. Silicon dioxide + Hydrofluoric Acid Silicon Tetrafluoride + water

SiO2 + 4 HF → SiF4 +2 H2O

12. Sulfuric Acid Sulfur Trioxide + Water H2SO4 → SO3 + H20 BalancedNotes #5: Reaction Types

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COMBINATION (synthesis) Reactions

Two or more reactants form a SINGLE product. Use crisscross rules to form compounds.

1. 4 Al (s) + 3 O2 (g) 2 Al2O3

2. 2 S (s) + 3 O2 (g) 2 SO3

3. 2 Cu (s) + _____ S (s) _____ Cu2S *_____ Cu (s) + _____ S (s) _____ CuS *

4. 2 Fe (s) + _____ O2 (g) 2 FeO *5. 4 Fe (s) + 3 O2 (g) 2 Fe2O3 *

*Elements with multiple charges can form two different compounds depending if limited or in excess.

Reactions between WATER and NONMETAL OXIDES usually give an ACID. (H + polyatomic)

6. _____ SO3 (g) + _____H2O (l) _____H2SO4 Balanced

7. _____ N2O5 (g) + _____H2O (l) 2 HNO3

Reactions between WATER and METAL OXIDES usually give a BASE. (Metal + OH)

8. _____ CaO (s) + _____H2O (l) _____Ca(OH)2 Balanced

9. _____ Na2O (s) + _____H2O (l) 2 Na(OH)

10. _____MgO (s) + _____H2O (l) _____Mg(OH)2 Balanced

DECOMPOSITION Reactions

Single reactant broken down into TWO OR MORE simpler products. 1. CaCO3 (s) CaO + CO2

2. 2 H2O (l) + electricity 2 H2 + O2

3. 2 Ag2O (s) 4 Ag + O2

4. H2O2 (l) H2 + O2

Metal Carbonate decomposes to METAL OXIDE and CO2.

5. Nickel (II) carbonate Nickel(II)oxide + _____

NiCO3 → NiO + CO2 Balanced

Metal Hydroxide decomposes to METAL OXIDE and H2O .

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6. Calcium hydroxide Calcium Oxide and Water

Ca(OH)2 → CaO + H2O Balanced

Metal Chlorate decomposes to METAL CHLORIDE and O2.

7. Potassium chlorate (s) Potassium Chloride + oxygen gas

2 KClO3 → 2 KCl + 3 O2

Some Acids decompose to NONMETAL OXIDE & H2O.

8. Carbonic Acid → ________________________________ + H2O

H2CO3 → CO2 + H2O Balanced

SINGLE REPLACEMENT Reactions

An element in reactant replaces a* similar element.o A + BC B + AC A = Metalo D + BC C + BD D = Nonmetal

1. _____ Zn (s) +_____ H2SO4 (aq) _____ ZnSO4 + _____ H2

2. _____ K (s) + 2 H2O (l) _____ KO2 + 2 H2

3. _____ Sn (s) + _____ NaNO3 (aq) No Reaction4. _____ Cl2 (g) + 2 NaBr (aq) 2 NaCl + _____ Br2

5. 2 Al (s) + 3 H2SO4 (aq) _____ Al2(SO4)3 + 3 H2

6. _____ KCl +_____ I2 No Reaction

7. Would copper replace chromium in a reaction? No

8. Would calcium replace mercury in a reaction? YES

9. What element would platinum be able to replace in a reaction.Gold

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You can use the activity series to predict whether or not certain reactions will occur. A specific metal can replace any metal listed below it that is in a compound. It cannot replace any metal listed above it. For example, copper atoms replace silver atoms in a solution of silver nitrate. However, if you place a silver wire in aqueous copper (II) nitrate, the silver atoms will not replace the copper. Silver is listed below copper in the activity series, so no reaction occurs. The letters NR (no reaction) are commonly used to indicate that a reaction will not occur.

**This applies to the nonmetal halogens as well. They increase as you move up the family.

Activity SeriesMetal

LithiumPotassium

BariumCalciumSodium

MagnesiumAluminumManganese

ZincChromium

IronCobaltNickel

TinLead

(Hydrogen)Copper

MercurySilver

PlatinumGold

DOUBLE REPLACEMENT Reactions

Exchange of positive ions between two compounds.o A+B- + C+D- A+D- + C+B-

o Products are usually a precipitate, gas or molecular compound such as water.

Formation of a gas:

1. _____ FeS (s) + 2 HCl (aq) _____ FeCl2 +_____ H2S

Acid-Base Reactionso When acids and bases react, they generally produce Salt and water. The common salt used to flavor foods

such as french fries or scrambled eggs is sodium chloride, NaCl, and is a product of some acid-base reactions. However, there are many other salts used in chemistry that are produced from acid-base reactions.

o Examples: HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l) acid base salt water

H2SO4 (aq) + 2KOH (aq) → K2SO4 (aq) + 2H2O (l) acid base salt water

2. HCl (aq) + Ca(OH)2 (aq) → CaCl2 (aq) + H2O (l)

3. _____ H2SO4 (aq) + 2 NaOH (aq) _____ Na2SO4 +2 H2O

4. 2 HCl (aq) + _____ Mg(OH)2 (aq) _____ MgCl2 + 2 H2O

Precipitation Reactions: When two solutions of ionic compounds are mixed, a product that can form is an insoluble salt called a precipitate. The precipitate is a solid that falls out of solution and is indicated as a solid on the product side of a chemical equation. We use a solubility chart to determine what compounds will combine to form a precipitate.

o Example: AgNO3 (aq) + NaCl (aq) → AgCl (s) + NaNO3 (aq) Precipitate

5. _____ Na2SO4(aq) + _____ Ba(NO3)2(aq) → _____ BaSO4(s) + 2 NaNO3(aq)

6. _____ K2CO3 (aq) + _____ Sr(NO3)2 (aq) → _____ SrCO3 (s) + 2 KNO3 (aq)

7. _____ BaCl2 (aq) _____ + K2CO3 (aq) _____ BaCO3 + 2 KCl

8. _____ SrBr2 (aq)_____ + (NH4)2 CO3 (aq)_____ SrCO3 +2 NH4Br

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COMBUSTION Reactions

Oxygen reacts with hydrocarbon to produce CARBON DIOXIDE and WATER for complete combustion. (Incomplete combustion results when NOT enough oxygen is available)

1. 2 C6H6 + 15 O2 12 CO2 + 6 H2O

2. 2 C8H18 + 25 O2 16 CO2 + 18 H2O

3. _____ C6H12O6 + 6 O2 6 CO2 + 6 H2O

4. 2 C10H22 + 31 O2 20 CO2 + 22 H2O

OXIDATION-REDUCTION (Redox) Reactions

Electrons are transferred from one atom to another through the change in oxidation numbers.o Oxidation Reduction

Loss of e- Gain of e-

Rules: o Assign Charges (oxidation #) to all atoms in equation.o Identify atoms that are oxidized & atoms reduced.o Determine the change in charge (oxidation #) for atoms that are oxidized or reduced.o Draw line connecting atoms oxidized or reduced & write the net change in oxidation #.

Oxidation numbers are a convenient way of identifying redox reactions and also indicating which element is oxidized and which is reduced. Here's an example - the reaction between sodium metal and chlorine gas:

2 Na + Cl2 → 2 NaCl

It is often useful to write the oxidation number for every element, in every compound, above the element in the equation. Thus for our reaction we have:

0 0 +1 -1

2 Na + Cl2 → 2 NaCl

o Be sure to note that the balancing coefficients in the equation (the "2" in front of Na and in front of NaCl) do not affect the value of the oxidation numbers. We'll return to these coefficients soon.

Since no electrons are transferred in acid-base reactions or precipitation reactions, acid-base reactions and precipitation reactions are not redox reactions.

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*OIL RIG = ox is loss, red is gain

A chart is a useful way for us to summarize the changes in oxidation number for each element:

element Initial ox. # Final ox # change in electrons (e-) oxidized or reduced

Na 0 → +1 lost 1 e- oxidized

Cl 0 → -1 gain 1 e- reduced

We see several important things in our table -

Since oxidation numbers did change, this was a redox reaction Na's oxidation number increased - from 0 on the reactant side to +1 on the product side. An element

becomes more positive by losing electrons. **Loss of electrons is Oxidation (LEO) Cl's oxidation number decreased, from 0 to -1, as chlorine gained electrons. **Gain of electrons is

Reduction (GER)

Practice: For the following redox reactions, write what elements are oxidized and reduced:

1. 2 Fe + 3 Cl2 2 FeCl3


Fe 0 → +1 lost 1 e- oxidized

Cl 0 → -1 gain 1 e- reduced

2. 2 KBr + Cl2 2 KCl + Br2


K +1 → +1 No change No change

Br -1 → 0 Lost 1 e- Oxidized

Cl 0 → -1 Gained 1 e- Reduced

3. N2 (g) + 3 H2 (g) 2 NH3 (g)


N 0 → -3 Gained 3 e- Reduced

H 0 → +1 Lost 1 e- oxidized

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4. 2 Mg (s) + O2 (g) 2 MgO (s)


Mg 0 → 2+ Lost 2 e- Oxidized

O 0 → 2- Gained 2 e- reduced

Reducing agent: loss = oxidizedOxidizing agent: gain = reduced

Academic Chemistry Unit 7 Test Review

1. How is the modern periodic table arranged?

2. Rows on the periodic table are called . Columns are called .

3. Be able to identify Metals, Nonmetals, and Metalloids on the periodic table.

4. Know the names for groups 1, 2, 3-12, 17 and 18.

5. Atoms across a period are arranged according to .

6. What are two ways in which atoms within a group are similar to one another?

7. How did Mendeleev arrange the periodic table?

8. How did Mosley arrange the periodic table?

9. Name 3 elements in the same period as nitrogen:

10. Name 3 elements in the same group as oxygen:

11. What is an anion?

12. Are anions larger or smaller than their neutral atoms?

13. What is a cation?

14. Are cations larger or smaller than their neutral atoms?

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15. What is the octet rule?

16. Elements in group 2A tend to form ions with what charge?

17. Elements in group 7A tend to form ions with what charge?

Describe the following trends across a period and down a group:ElectronegativityDefinition:

Explain Reason for Trend Label & Describe the trends

First Ionization EnergyDefinition:


Atomic RadiiDefinition:


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Ionic RadiiDefinition:


18. ____ NaBr + ____ H3PO4 ____ Na3PO4 + ____ HBr Type of rxn: ______________________

19. ____ C2H4 + ____ O2 ____ CO2 + ____ H2O Type of rxn: ______________________

20. ____ Mg + ____ Fe2O3 ____ Fe + ____ MgO Type of rxn: ______________________

21. ____ PbSO4 ____ PbSO3 + ____ O2 Type of rxn: ______________________

22. ____ H2SO4 + ____ NH4OH ____ H2O + ____ (NH4)2SO4 Type of rxn: ______________________

23. __HCl (aq) + __NaOH (aq) → __NaCl (aq) + __H2O (l) Type of rxn: ______________________

24. __Pb(NO3)2 (aq) + __H2SO4 (aq) → __PbSO4 (s) + __HNO3 (aq) Type of rxn: ______________________

25. __HNO3 (aq) + __KOH (aq) → __KNO3 (aq) + __H2O (l) Type of rxn: ______________________

26. __Al (s) + __HCl (aq) → __AlCl3 (aq) + __H2 (g) Type of rxn: ______________________

27. __I2O5 (s) + __CO (g) → __I2 (s) + __CO2 (g) Type of rxn: ______________________

28. __H2SO4 (aq) + __KOH (aq) → __H2O (l) + __K2SO4 (aq) Type of rxn: ______________________

29. __Cu(s) + __HNO3(aq) → __Cu(NO3)2(aq) + __NO2(g) + __H2O(l) Type of rxn: ______________________

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30. __Na3PO4 (aq) + __FeCl3 (aq) → __NaCl (aq) + __FePO4 (s) Type of rxn: ______________________

31. __Pb(C2H3O2)2 (aq) + __HCl (aq) → __PbCl2 (s) + __HC2H3O2 (aq) Type of rxn: ______________________

32. __Bi(OH)3(s) + __Na2SnO2(aq) → __Bi(s) + __Na2SnO3(aq) + __H2O(l) Type of rxn: ______________________

33. __HNO3 (aq) + __Mg(OH)2 (aq) → __H2O (l) + __Mg(NO3)2 (aq) Type of rxn: ______________________

34. __Cl2 (g) + __H2O (l) → __HCl (aq) + __HClO (aq) Type of rxn: ______________________

35. __Na3PO4 (aq) + __Pb(NO3)2 (aq) → __NaNO3 (aq) + __Pb3(PO4)2 (s) Type of rxn: ______________________

36. __NaOH (aq) + __CaCl2 (aq) → __NaCl (aq) + __Ca(OH)2 (s) Type of rxn: ______________________

37. __Na2CO3 (aq) + __Ca(NO3)2 (aq) → __CaCO3 (s) + __NaNO3 (aq) Type of rxn: ______________________

38. __S (s) + __HNO3 (aq) → __SO2 (g) + __NO (g) + __H2O(l) Type of rxn: ______________________

39. __AgNO3 (aq) + __H2S (aq) → __Ag2S (s) + __HNO3 (aq) Type of rxn: ______________________

40. __HCl (aq) + __LiOH (aq) → __H2O (l) + __LiCl (aq) Type of rxn: ______________________

41. __KMnO4(aq) + __HCl(aq) → __MnCl2(aq) + __Cl2(g) + __H2O(l) + KCl(aq) Type of rxn: ______________________

42. __(NH4)2S (aq) + __Co(NO3)2 (aq) → __CoS (s) + __NH4NO3 (aq) Type of rxn: ______________________

43. __H3PO4 (aq) + __Ca(OH)2 (aq) → __H2O (l) + __Ca3(PO4)2 (aq) Type of rxn: ______________________

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