chemical reactions chapter 7 a way to describe what happens in a chemical reaction. 1)tells us what...

Chemical Reactions

Chapter 7

• A way to describe what happens in a chemical reaction.

1) Tells us what substances are involved with the reaction

2) Tells us how much of each substance is needed in a chemical reaction

Reactants & Products

• REACTANTS• The substances

that go into a chemical equation.

• Left side of the arrow

• PRODUCTS• The substances

produced or made in a chemical reaction

• Right side of the arrow

What do the numbers mean?

• Coefficients : the numbers placed in front of an element or compound. Tells us how many molecules or units of that substance

are present. 3 NaOH• Subscripts : the small numbers placed after

and below an element or ion. Tells us how

many atoms are present. CaCl2

More chemical equations:

• Magnesium + oxygen magnesium oxide

• Mg + O2 MgO

• What is wrong with this equation?

Balancing Chemical Equations

1. Can only add or change coefficients NEVER subscripts.

2. Balance hydrogen atoms last.

3. Balance oxygen atoms second to last.

Law of Conservation of Mass

• Matter can neither be created nor destroyed.

• Developed by Antoine Lavoisier (1743-1794)

The Mole• A special unit in

chemistry used to measure the number of particles in a specific amount of mass.

• One mole equals exactly 6.02 x 1023 atoms. (Avogadro’s number)

A mole is used the same way as:

Working problems with moles

• 1 mole of iron would have a mass of how many grams?

• 64 grams of sulfur is equal to how many moles?

• 1/2 or 0.5 moles of water has a mass of ?

1. Synthesis reactions

• A type of chemical reaction where two or more reactants combine to form a single substance.

• General equation:

• A + B AB

Examples of synthesis reactions:

• Formation of salt from chlorine gas and solid sodium.

• Magnesium oxide formation.

• Rusting of metals

2. Decomposition reactions

• When a complex substance breaks down into two or more simpler substances.

• Generic equation:

• AB A + B

Decomposition examples

• Electrolysis of water into hydrogen gas and oxygen gas.

• Production of cement.

• Sodium azide in safety air bags. Page 201

Single-Replacement reactions

• When an uncombined element replaces another element in compound.

• Generic equation:

• A + BC AC + B• Which of the two combined elements are

replaced?

• “Like” replaces “like”

• Metals replace metals and non-metals replace non-metals.

• Examples:

• Zinc + copper sulfate

• Calcium + water

• Fluorine gas + potassium bromide

4. Double Replacement

• Also known as ionic exchange reactions.

• Generic equation:

• AB + CD AD + CB• Example:

– Lead nitrate + potassium iodide

5. Combustion reactions

• When a substance rapidly reacts with oxygen to produce light and/or heat.

• “burning”

• Example:

• Bunsen burner: methane + oxygen

• Sometimes these reactions are classified as synthesis reactions.

REDOX

• OXIDATION• When a metal

combines with oxygen.

• Rusting Any process where an

element loses electrons

• REDUCTION• When an element

gains electrons during a reaction.

Energy changes in chemical reactions

• 2 things that always change during a chemical reaction:

1. The properties of the reactants.

2. The amount of energy present.• Energy change occurs because of breaking

and formation of chemical bonds.

Exothermic reactions

• A chemical reaction that releases energy to its surroundings.

• Often in the form of heat. Reaction feels warm or hot.

• Examples:• Mg + HCL

Endothermic reactions

• A chemical reaction that absorbs energy from its surroundings.

• Often feel cool.• Examples:• Ice-packs

Energy diagrams

• Exothermic • Endothermic

Reaction rates

• Defined as how quickly a reaction occurs.

• Collision theory = The more collisions that occur at the atomic level the faster the reaction will go

Factors that affect rate of reaction

1) Temperature

2) Surface Area

3) Stirring

4) Concentration

5) Catalysts

TEMPERATURE

• Generally, the higher the temperature of the reactants the faster the rate of reaction.

SURFACE AREA

• The greater the surface area, the faster the rate of reaction.

• Example:

STIRRING

• Speeds up the molecules thus increasing the number of collisions

CONCENTRATION

• Defined as the number of particles in a given unit of volume.

• Example: • 12 M HCl• 6 M HCl• 1 M HCl

CATALYSTS

• A substance that speeds up a reaction without being used up or directly involved in the reaction.