Chemical reactions
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Chemical Reactions. Chapter 4-1 – 4-4; 4-7 – 4-12 (and a little of Chapter 3). Key concepts. Identify metals, non-metals, and metalloids on the periodic table. Identify key characteristics of metals and non-metals. Introduction to strong acids/bases and weak acids/bases in solution.

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Chemical Reactions

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Chemical reactions

Chemical Reactions

Chapter 4-1 – 4-4; 4-7 – 4-12

(and a little of Chapter 3)


Key concepts

Key concepts

  • Identify metals, non-metals, and metalloids on the periodic table.

  • Identify key characteristics of metals and non-metals.

  • Introduction to strong acids/bases and weak acids/bases in solution.

  • Write formula unit, total ionic, and net ionic equations for aqueous solutions.

  • Learn how to assign oxidation numbers to atoms.

  • Learn how to identify combination reactions, decomposition reactions, displacement reactions, and metathesis reactions.


Metals and nonmetals

Metals and nonmetals

  • We can classify elements according to atomic structure (current periodic table).

  • Mendeleev (1872) came up with essentially the same table before quantum mechanics developed, using general chemical properties of different elements.

  •  there must be some relationship between atomic structure and chemical properties.


Metals and non metals physical properties

Metals and non-metals—physical properties


Metals and non metals chemical properties

Metals and non-metals—chemical properties


Common groups in periodic table

Common groups in periodic table

  • IA: Alkali Metals

  • IIA: Alkaline earths

  • VIA: Chalcogens

  • VIIA: Halogens

  • VIIIA: Noble gases (rare gases)


Aqueous solutions

Aqueous Solutions

  • Many of the reactions we encounter in nature occur between substances in a solution of water. We call these aqueous solutions, and introduce them here

  • We’ll talk specifically about the process of solvation (dissolution) later in the semester.


Solutions

Solutions

  • A solution is a ____________________.

  • A Solution contains a solute and a solvent.

  • Solute:

  • Solvent:

  • There are a variety of ways to express the concentration of solute in a solvent.


By mass

% by mass

This is mass of the solution,

not mass of the solvent.


Chemical reactions

  • In very dilute solutions, % solute is an awkward unit to use. Often, parts-per-million (ppm) or parts-per-billion (ppb) are used instead.


Molarity

Molarity

  • Chemists will often use molarity because it is more directly related to the moles of solute (and becomes more useful in analyzing chemical reactions).

Not volume of solvent.


Solution dilution

Solution dilution…..

  • Suppose we change the volume of a solution, but we leave the amount of solute the same….


Solutions in chemical reactions

Solutions in chemical reactions

  • Just as we dealt with the mass of reactants and products in a chemical reaction by converting to moles, we may also use the molarity of a solution in analyzing chemical reactivity.

  • The key is to always compare moles to moles.


Electrolytes

Electrolytes

  • Electrolytes are substances whose aqueous solutions __________________.

  • Strong electrolytes:

  • Weak electrolytes:

  • Non-electrolytes:

  • Electrolyte solutions carry electric current through ___________________.

  • Demo: NaCl, acetic acid, sugar conductivity


Dissociation and ionization

Dissociation and ionization

  • Both dissociation and ionization processes form ionic solutions.

  • Dissociation forms solutions when _______________________________.

  • Ionization forms solutions when ________________________________.


Acids and bases

Acids and bases

  • We’ll talk more about acids and bases in Chapters 10-11. As an introduction,

  • Acids are substances that produce _______ in aqueous solutions.

  • Bases are substances that produce ______ in aqueous solutions.


Ionic salts

Ionic salts

  • An ionic salt is ____________ . Salts form during reactions of _________________.

  • Strong electrolyte solutions are made from three types of solutes:

    • Strong acids

    • Strong bases

    • Most soluble salts

  • Why do these compounds form strong electrolyte solutions?


Strong acids and weak acids

Strong acids and weak acids

  • Strong acids ionize _________ in aqueous solutions.

    • (HCl dissolved in water)

    • There are relatively few strong acids. KNOW the acids listed in Table 4-5 in text.

  • Weak acids ionize slightly in dilute aqueous solutions, and are in ________ with their conjugate anions.

  • There are a multitude of weak acids, many play roles in biological processes.

  • The ionization/dissociation of a strong acid is an ___________ reaction.

  • The ionization of a weak acid is a _____________ reaction.


Strong bases and weak bases

Strong bases and weak bases

  • The behavior of strong bases and weak bases is similar to strong acids and weak acids. Can you predict the type of reactions you will see with each?

  • Insoluble bases: contain ionic hydroxides, but do not produce strong basic solutions because ____________.


Solubility rules

Solubility rules

  • Solubility rules (like Lewis structure rules) are not absolute. They provide general guidelines on what is soluble and what is not.

  • Solubility rules help us predict what products may be observed in a given chemical reaction.

    • Will a precipitate form during a reaction? What will it be?


What is soluble

What is soluble?

  • Soluble:

    • Common inorganic acids and low-molecular-weight organic acids

    • Common compounds containing IA element ions and NH4+

    • Common nitrates, acetates, chlorates, perchlorates

    • Common chlorides, except AgCl, Hg2Cl2, PbCl2

    • Common bromides and iodides similar to chlorides, but with more exceptions

    • Common fluorides except those with Mg2+, Ca2+, Sr2+, Ba2+, and Pb2+

    • Common sulfates, except those with Pb2+, Ba2+, and Hg2+. (CaSO4, SrSO4, Ag2SO4 moderately sol.)


What is insoluble

What is insoluble?

  • Insoluble:

    • Common metal hydroxides, except those with IA metal cations and IIA metal cations Ca2+ and heavier.

    • Common carbonates, phosphates, and arsenates, except those with IA metal cations and NH4+. (MgCO3 mod. Sol.)

    • Common sulfides, except IA and IIA metal cations and NH4+.

  • You should know the solubility rules outlined on p 134 and Table 4-5 in your text.

    Let’s look at some examples…


Aqueous solution reactions

Aqueous solution reactions

  • For aqueous solutions, we can manipulate the way a chemical reaction is written to illustrate the processes that are taking place in a solution.

  • The different types of equations include…


1 formula unit equations

1. Formula unit equations

  • Formula unit equations show __________ for all compounds in the solution.

  • Example: amalgamation of mercury: mercury (II) nitrate combines with copper metal to produce mercury metal and copper (II) nitrate.

  • Write out the formula unit equation….


2 total ionic equation

2. Total ionic equation

  • Formula shows the form of the substance when dissolved in aqueous solution….

  • How would we write the total ionic equation for the reaction in #1?


3 net ionic equation

3. Net ionic equation

  • Net ionic equations include only those ions involved in the reactions. The ions not participating in the reaction (called _______ ______) are not included.

  • What is the net ionic equation for our reaction?

  • (make sure charges on both sides balance)

  • (Check flowchart on p 137 as a guide for writing equations)


Oxidation numbers

Oxidation numbers

  • Oxidation numbers are bookkeeping tools that help us keep track of the transfer of electrons during a chemical reaction. They are essential in balancing types of reactions known as reduction-oxidation reactions, or redox reactions. Redox reactions will be discussed in more detail in Chemistry 106.

  • Oxidation numbers are also useful in checking the validity of a written chemical formula.


Rules for oxidation numbers

Rules for oxidation numbers

  • NUMBER ONE RULE: THE SUM OF ALL OXIDATION NUMBERS IN AN ION, ATOM, OR MOLECULE MUST EQUAL THE CHARGE OF THE SPECIES.

    • So, for instance, the sum of all oxidation numbers for the atoms in a neutral molecule must equal the charge, i.e., they must equal zero.

    • The sum of oxidation numbers on an ion must equal the charge of that ion


Other rules

Other rules

  • In an elemental compound, all atoms are 0.

  • Monoatomic cations have an ox. number equal to their charge.

  • Oxygen is always –2, except when in peroxides (when it is –1), superoxides (-1/2), or when bonded to F.

  • Hydrogen is always +1, except when bonded to a metal in a metal hydride, (like NaH), when it is –1.


Chemical reactions

  • Generally, elements closer to fluorine have the greater oxidation number in a compound.

  • There are several other oxidation number rules given on p 138-139, and you should be familiar with them. But, these rules build on the 5 main rules given above, and on the common charges for ions.

  • More examples:


Types of chemical reactions

Types of chemical reactions

  • (the text goes through several examples of these types of reactions. You should make sure you understand what is going on in these examples. We will not cover all the examples in class, but you will be responsible for understanding the different classes of reactions)


I redox reactions

I. Redox reactions

  • How to identify:

  • Oxidation is indicated by ______________

  • Reduction is indicated by ______________

  • The substance that is oxidized is also the _________ __________.

  • The substance that is reduced is also the __________ __________.

  • Many of the other types of reactions we will discuss are also redox reactions

  • Examples:


Combination reactions

Combination reactions

  • Two or more reactants combine to form one product.


Examples of combination reactions

Examples of combination reactions

  • Rust

    • What is the reaction?

    • Is this also a redox reaction? How can you tell?

  • Water formation

    • What’s the reaction?

    • Is it redox?

  • CaO (s) + H2O(l) → Ca(OH)2 (aq)

    • Redox?


Decomposition reactions

Decomposition reactions

  • One reactant decomposes to form more than one product.


Examples of decomposition reactions

Examples of decomposition reactions

  • Electrolysis of water

    • Equal moles of gas produce equal volumes

    • What is observed here?

  • Potassium chlorate decomposition

  • Thermal decomposition


Displacement reactions

Displacement reactions

  • One element displaces another one from a compound.

  • Displacement reactions are always redox reactions.


Activity series of the elements

Activity series of the elements

  • More active metals displace less active metals (or hydrogen) from aqueous solution. The oxidized form of the more active metal and the reduced form of the less active metal are products.

  • Examples:

    • Amalgamation of mercury on Cu (done before)

    • Displacement of Zn in acidic solution

    • Displacement of H from water


Metathesis reactions

Metathesis reactions

  • Two ions “switch partners” to form new ionic compounds, with no change in oxidation number.

  • Two main types:

    • acid-base neutralization

    • Precipitation reaction


Acid base neutralization

Acid base neutralization

  • Acid and base react to form non-electrolyte product. Many times, the base is a ______________ compound, and the products are a __________ and __________.


Precipitation reactions

Precipitation reactions

  • One of the products is an insoluble salt (important to remember solubility rules), and will precipitate from (come out of) solution.

  • Examples:

    • Silver nitrate and sodium chloride

    • Sodium carbonate and calcium chloride


Summary

Summary

  • Table 4-16 in text helps summarize the different types of reactions….use this to help review.


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