HOW FAR? HOW FAST?

IGCSE Class

Mr. Markus

Chemical Properties

• Chemical properties are those that can be observed only when there is a change in the composition of the substance.

• Inability to react is also a chemical property.

Chemical Change:

The change of one or more substances

into other substances.

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Other examples of chemical changes:

• Decomposition
• Explosion
• Oxidation
• Corrosion
• Tarnish
• Fermentation
• Burning (combustion)
• Rotting

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Recognizing a Chemical Change�

• energy exchange
• production of a gas
• color change
• formation of a precipitate

• Another term for chemical change is chemical reaction.

Chemical Reactions and Energy

• All chemical changes involve an energy change.

Energy

• Energy is either taken in or given off as the chemical change takes place.
• Energy is the capacity to do work.
• Work is done when a force acts to move an object.

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• change involves a rearrangement of atoms
• A chemical reaction involve the breaking of chemical bonds in the reactants and the formation of chemical bonds in the product.

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ENERGY CHANGES IN REACTION

Exothermic Reactions:

• Reactions that give off heat energy.

Example:

Burning wood

Endothermic Reactions:

• Reactions that absorb heat energy.

Example:

decomposition of

water

Photosynthesis

• Photosynthesis is an endothermic process.

• Green plants, algae, and many kinds of bacteria carry out photosynthesis.

Exothermic Reactions

• Releases energy to

its environment.

– The energy released

as products form is

greater than the

energy required to

break the bonds in

the reactants.

Endothermic Reaction

• A chemical reaction

that absorbs energy

from its

surroundings.

– More energy is

required to break the

bonds in the

reactants than is

released by the

formation of

products.

• Breaking bonds requires energy. It is endothermic.

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• Making new bonds gives out energy. It is exothermic.

Bond energy

When a bond is broken, energy is absorbed. When a bond is formed, energy is released. 

Example

Cl2 + 58 Kcal Cl + Cl    (endothermic, ΔH is positive)

Cl + Cl Cl₂ + 58 Kcal    (exothermic, ΔH is negative)

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Notice that in the potential diagram for the above reactions the molecule Cl₂ is more stable than the 2 atoms of Cl. This is why chlorine can be found in nature as a diatomic molecule.

Example

H₂(g), N₂(g), O₂(g), F₂(g), Cl₂(g), Br₂(l), and I₂(s)

Burning of Methane�CH₄ +2O₂🡪 CO₂ + 2H₂O

Ex: Burning of Methane�CH₄ + 2O₂🡪 CO₂ + 2H₂O

CH₄(g) + 2O₂(g) 🡪 CO₂(g) + 2H₂O(g)

The left side involves bond breaking and energy needs:

Four C – H bonds 4 x 435 kJ/mol = 1740 kJ/mol

Two O = O bonds 2 x 497 kJ/mol = 994 kJ/mol

total energy needed = 2734 kJ/mol

The right side involves bond making and gives out energy:

Two C = O bonds 2 x 803 kJ/mol = 1606 kJ/mol

Four O – H bonds 4 x 464 kJ/mol = 1856 kJ/mol

total energy given out = 3462 kJ/mol

The heat of reaction, ΔH is the energy change on going from reactants to products

ΔH = (energy needed to break bonds) – (energy given out when bonds form)

= 2734 – 3462

= -728 kJ/mol

Example: Composition of Water�2H₂ + O₂🡪 2 H₂O

The Law of Conservation of Mass �(Antoine Lavoisier)

• In any chemical or physical change, mass is neither created or destroyed

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• Mass is CONSTANT

Example:

• 4g of hydrogen react with 32 g of oxygen to form water.
• How many grams of water are formed?

Answer: 36g