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CHEMISTRY - CLUTCH

CH.3 - CHEMICAL REACTIONS

CONCEPT: ALLOYS

An alloy represents a _________________ mixture composed of two or more elements in which at least one of the elements is a metal.

• The predominant metal component, which is can be up to 90% or higher in terms of composition is referred to as the ____________ or ____________ metal.

• The minor components, which usually average around 1% in terms of composition, are referred to as the _______________________.

Alloys can be created by three different methods:

1) Heating the alloy components into liquids, mix them together and allowing them to cool into a ____________________.

2) ____________________ : turning the components into powders, mixing them together and allowing them to fuse.

3) ____________________: firing beams of ions at the surface of the host metal and allowing other components to mix.

Alloys are classified into two major types:

• A(n) _______________________ alloy is where some of the host metal atoms have been replaced by other metal atoms that have a similar size.

• A(n) _______________________ alloy is where the empty spaces between the host metal atoms have been taken up by smaller metal atoms.

Other common alloys that are good to remember include:

______________________ – copper (host metal), tin, manganese, phosphorus, aluminum, silicon.

______________________ – tin (host metal), copper, lead, antimony.

______________________ – iron (host metal), chromium with very small amounts of carbon, nickel, manganese & molybdenum.

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CONCEPT: COMMON POLYATOMIC IONS

Polyatomic ions are compounds made up of different elements, usually only ____________, and possess a ____________.

Singly Charged Cation (Positive Ion)

NH4+

Ammonium

Doubly Charged Cation (Positive Ion)

Hg22+

Mercury (I)

Singly Charged Anions (Negative Ions)

CH3CO2– or C2H3O2

–

Acetate

CN–

Cyanide

OH–

Hydroxide

MnO4–

Permanganate

NO3–

Nitrate

Nitrite

Doubly & Singly Charged Anions (Negative Ions)

HPO42–

Hydrogen Phosphate

H2PO4–

Dihydrogen Phosphate

HCO3–

Hydrogen Carbonate or Bicarbonate

HSO4–

Hydrogen Sulfate or Bisulfate

Doubly Charged Anions (Negative Ions)

CO32–

Carbonate

CrO42–

Chromate

Cr2O72–

Dichromate

O22–

Peroxide

SO42–

Sulfate

Sulfite

Triply Charged Anions (Negative Ions)

PO43–

Phosphate

Phosphite

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CONCEPT: POLYATOMIC IONS w/ HALOGENS

Polyatomic ions containing halogens are sometimes referred to as __________ halogens or halogen ________________.

These compounds share 4 common characteristics:

1.

2.

3.

4.

These compounds use the same system for naming:

PRACTICE: Name each of the following compounds.

a. BrO4 – b. FO2 –

c. ClO – d. IO3 –

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CONCEPT: NAMING MOLECULAR COMPOUNDS Features: _________________ & _________________

Because molecular compounds combine in different proportions to form different compounds, we must use numerical

prefixes.

Rules for Naming: a. The first nonmetal is named normally and uses all numerical prefixes except ___________________. b. The second nonmetal keeps its base name but has its ending changed to _____________________. EXAMPLE: Write the formula for each of the following compounds.

a. Disulfur monobromide b. Iodine Tetrachloride

PRACTICE: Give the systematic name for each of the following compounds:

a. CO b. N2S4 c. IO5

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CONCEPT: IONIC COMPOUNDS

In the early days of chemistry, newly discovered compounds were given fancy names such as morphine, quicklime and

muriatic acid. Since then thousands of new compounds have been discovered and named under a system called

_____________________________.

Metals tend to __________ electrons to become positively charged ions called _______________.

Nonmetals tend to __________ electrons to become negatively charged ions called _______________.

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CONCEPT: NAMING BINARY IONIC COMPOUNDS Features: ___________________ & ___________________

Rules for Naming: a. The metal is named and written first.

• If the metal is a transition metal we must use a _________________________ to describe its positive charge.

b. The nonmetal keeps its base name but has its ending changed to ___________________.

EXAMPLE: Provide the molecular formula or name for each of the following compounds.

a. Calcium phosphide b. CoO

PRACTICE: Provide the molecular formula or name for each of the following compounds.

a. AlBr3 b. Lead (IV) sulfide c. SnO2

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CONCEPT: NAMING IONIC COMPOUNDS w/ POLYATOMICS Features: _________________ & _________________

Rules for Naming:

a) The metal keeps its name and is named and written first.

• If the metal is a transition metal we must use a _____________________ to describe its positive charge.

b) Name the polyatomic as you would normally.

EXAMPLE: Write the formula for each of the following compounds:

a. Iron (III) Acetate b. Copper (I) phosphate

c. Strontium Carbonate d. Ammonium Nitrite

EXAMPLE: Give the systematic name for each of the following compounds:

a. Pb(CrO4)2 b. Ga(ClO4)3

c. Mn(HSO4)2 d. Ba(CN)2

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CONCEPT: NAMING IONIC HYDRATES Features: _________________ & _________________

CuSO4 5 H2O

Rules for Naming the Ionic Compound portion: a. The metal is named normally and written first.

• If the metal is a transition metal we must use a ________________________ to describe its positive charge.

b. The nonmetal keeps the first part of its name but has its ending changed to ___________________.

c. Name the polyatomic as you would normally.

Rules for Naming the H2O portion:

a. The H2O portion will be called ___________________ .

b. To describe the number of H2O molecules use these prefixes.

EXAMPLE: Write the formula for each of the following compounds.

a. Calcium carbonate hexahydrate

b. Lead (IV) Sulfate pentahydrate


a) K2Cr2O7 · 3 H2O b) Sn(SO3)2 · 4 H2O

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CONCEPT: NAMING ACIDS

1. BINARY ACIDS Features: _______________________ + _______________________

Rules for Naming:

a. The prefix will be ___________________ .

b. Use the base name of the nonmetal.

c. The suffix will be ___________________ .

EXAMPLE: Write the formula for each of the following compounds:

a. Hydroiodic acid b. Hydroselenic acid c. Hydrofluoric acid


a. HBr b. H2S c. HCN

2. OXOACIDS or OXYACIDS Features: _______________________ + _______________________

Rules for Naming: a. If the polyatomic ion ends with –ate then change the ending to _____________________. b. If the polyatomic ion ends with –ite then change the ending to ______________________.

EXAMPLE: Give the systematic name or formula for each of the following compounds:

a. H2CO3 b. Nitric acid c. H2SO4

PRACTICE: Give the systematic name or formula for each of the following compounds:

a. Hypobromous acid b. HClO3 c. Acetic acid

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CONCEPT: EMPIRICAL FORMULA

The empirical formula is also known as the __________________________.

• It represents the _______________________________ ratio of moles

of each element in the compound.

The molecular formula is also known as the __________________________.

• It represents the _______________________________ ratio of moles of each element in the compound.

EXAMPLE: What is the empirical formula of dimethylhydrazine, C2H8N2, a colorless liquid used as a rocket fuel?

EXAMPLE: Elemental analysis of a sample of an ionic compound showed 2.82 g of Na, 4.35 g of Cl, and 7.83 g of O. What

is the empirical formula and name of the compound?

EXAMPLE: After a workout session, lactic acid (M = 90.08 g/mol) forms in muscle tissue and is responsible for muscle

soreness. Elemental analysis shows that this compound contains 40% C, 6.7% H and 53.3% O. Determine the molecular

formula.

C6H12O6 =

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CONCEPT: COMBUSTION ANALYSIS

Under a combustion reaction a compound made of _______________ or _______________ reacts with ______________ .

• The products formed will be ________________ and ________________ .

EXAMPLE: A 0.2500 g sample contains carbon, hydrogen and oxygen and undergoes complete combustion to produce

0.3664 g of CO2 and 0.1500 g of H2O. What is the empirical formula of the compound?

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CONCEPT: CHEMICAL COMPOSITION

The mass per mole of an element is called its _____________________________ (M).

The mass per mole of a compound is called its ____________________________ (M).

• They both have the units of _______________.

1. Elements. To find mass of an element just look up its atomic mass in the periodic table.

EXAMPLE: What is the total mass of each of the following elements?

a. Sodium b. Gold c. Mercury 2. Compounds. The mass of a compound is the sum of the individual masses of the elements in the chemical formula.

EXAMPLE: What is the total mass of each of the following compounds?

a. N2O5 b. C12H22O11 c. (NH4)3PO4

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CONCEPT: MASS PERCENT

Mass or weight percent is the percentage of a given element in a compound.

EXAMPLE: What is the percentage of carbon in sodium hydrogen carbonate, NaHCO3?

EXAMPLE: A sample of toothpaste contains tin (II), SnF2. Analysis of a 5.25 g sample contains 8.77 x 10-3 g of F. What is

the percentage of tin (II) fluoride in the sample?

PRACTICE: Hemoglobin contains 0.33% iron and has a molecular weight of 68 kg. How many iron atoms are in each

molecule of hemoglobin?

Mass Percent (%) =

MassComponent)(TotalMass)(

•100

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CONCEPT: BALANCING CHEMICAL REACTIONS

When balancing an equation always make sure the ________ and ________ of atoms on both sides of the arrow are equal.

EXAMPLE: Write balanced equations for each of the following by inserting the correct coefficients in the blanks:

a. ____ Al (s) + ____ Cl2 (g) ____ AlCl3 (s)

b. ____ Ba3(PO4)2 (s) + ____ KOH (aq) ____ K3PO4 (aq) + ____ Ba(OH)2 (aq)

c. ____ C4H10 (aq) + ____ O2 (g) ____ CO2 (g) + ____ H2O (l)

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CONCEPT: GROUP REACTIVITY

The central principle of Organic Chemistry is based on the _________________________________. • The reactivity of an organic compound is all based on which type is present.

C

CC

C

CC

CO

C

C

CC

CC

CH3

H3C

H3C

OH

Tetrahydrocannabinol (THC)

CH2CH2CH2CH2CH3

Alkane

Alkene

Alkyne

Alcohol

Amine

Aldehyde

Ketone

Carboxylic Acid

Ester

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CONCEPT: GROUP REACTIVITY (PRACTICE)

EXAMPLE: In each of the following molecules, identify the type(s) of functional groups present.

a. b.

O

OH

PRACTICE: In each of the following molecules, identify the type(s) of functional groups present.

a. b.

O

O

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CONCEPT: STOICHIOMETRIC REACTIONS

2 H2 (g) + 1 O2 (g) 2 H2O (g)

In the above equation the numbers that are in bold are called _______________________.

• They tell us the number of ______________ of each compound that reacts.

• This numerical relationship between compounds in a balanced equation is called __________________________.

STOICHIOMETRIC CHART

Before we get into solving stoichiometric reactions lets work out a plan of attack.

Entities means ______________________ , ______________________ or ______________________.

Entities of Given

Grams of Given

Entities of Unknown

Grams of Unknown

Moles of Given Moles of Unknown

Use this chart when given a chemical equation with the ____________ quantity of a compound or element and asked to find

the ____________ quantity of another compound or element.

EXAMPLE: How many grams of H2O are produced when 12.3 g H2 reacts?

2 H2 (g) + 1 O2 (g) 2 H2O (g)

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PRACTICE: STOICHIOMETRIC REACTIONS

EXAMPLE 1: The oxidation of chromium solid is represented by the following equation:

4 Cr (s) + 3 O2 (g) 2 Cr2O3 (s)

a. How many moles of chromium (III) oxide are produced when 34.69 g Cr reacts with excess oxygen gas?

b. How many grams of O2 were needed to produce 4.28 x 103 molecules Cr2O3?

EXAMPLE 2: If the density of ethanol, CH3CH2OH, is 0.789 g/mL, how many milliliters of ethanol are needed to produce 4.8

g of H2O in the following reaction?

CH3CH2OH (l) + 3 O2 2 CO2 (g) + 3 H2O (l)

PRACTICE: Dinitrogen monoxide gas decomposes to form nitrogen gas and oxygen gas. How many molecules of oxygen

are formed when 8.00 g of dinitrogen monoxide decomposes?

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CONCEPT: LIMITING REAGENT

In a chemical reaction the reactant that is consumed when a reaction occurs and determines the maximum amount of

product formed is called the _________________________________.

• The amount of product it forms is called the _________________________ yield.

The reactant that remains after the completion of the chemical reaction is called the ____________________ reactant.

EXAMPLE: Chromium (III) oxide reacts with hydrogen sulfide (H2S) gas to form chromium (III) sulfide and water:

Cr2O3 (s) + 3 H2S (g) Cr2S3 (s) + 3 H2O (l) [Balanced]

a. What is the mass of chromium (III) sulfide formed when 14.20 g Cr2O3 reacts with 12.80 g H2S?

b. Identify the limiting reactant, excess reactant and theoretical yield.

c. What mass of excess reactant remains?

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CONCEPT: PERCENT YIELD

The percent yield of a reaction is used to determine how effective the chemist was in creating their desired products.

• A high percent yield would signify that the reaction is _______________________________________.

EXAMPLE 1: A scientist performs an experiment in the laboratory and obtains 13.27 g Cr2S3. If his calculations on scratch

paper give him a theoretical yield of 18.23 g what is the percent yield?

PRACTICE: Consider the following balanced chemical reaction:

2 C6H6 (l) + 15 O2 (g) 12 CO2 (g) + 6 H2O (l)

a. If a 2.6 g sample of C6H6 reacted with excess O2 to produce 1.25 g of water, what is the percent yield of water?

b. If the above reaction only went to 75% completion, how many moles of CO2 would be produced if 1.57 x 10-5 molecules of C6H6 were reacted with excess oxygen?

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CHEMISTRY - CLUTCH CH.3 - CHEMICAL REACTIONSlightcat-files.s3.amazonaws.com/packets/admin_chemistry...CH.3 - CHEMICAL REACTIONS Page 2 CONCEPT: COMMON POLYATOMIC IONS Polyatomic ions