SCIENCE 10
UNIT A: ENERGY AND MATTER IN CHEMICAL CHANGE
Unit Focus Questions
- How has knowledge of the structure of matter led to other scientific advancements?
- How do elements combine? Can these combinations be classified and the products be predicted and quantified?
- Why do scientists classify chemical change, follow guidelines for nomenclature and represent chemical change with equations?
A3.0 Chemical change is a process that involves recombining atoms and energy flows
Key Concepts
In this section, you will learn about the following key concepts:- Chemical substances and human needs
- Evidence for chemical change
- Role and need for classification of chemical change
- Writing and balancing equations
- Law of conservation of mass
- The mole concept
Learning Outcomes
When you have completed this section, you will be able to:- Provide examples of household, commercial and industrial processes that use chemical reactions, and identify chemical reactions that are significant in our society.
- Describe evidence of chemical change
- Differentiate between endothermic and exothermic chemical reactions
- Translate word equations to balanced chemical equations and vice versa
- Classify chemical reactions into categories, including formation, decomposition, hydrocarbon combustion, single replacement and double replacement
- Starting with the reactants, predict the products of formation, decomposition, hydrocarbon combustion, single replacement and double replacement reactions.
- Define the mole and use Avogadro’s number (6.02x1023) to relate numbers of particles in a substance to the quantity of substance
- Interpret balanced chemical equations in terms of moles of chemical species, and relate the mole concept to the law of conservation of mass.
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A3.1 Important Examples of Chemical Change
Reactant – substance that reacts in a chemical reaction to form another substance or substancesProduct –new substance produced in a chemical reaction
Reactions That Form Gases
Reactions That Form Solids
Showing States in Chemical Formulas
Elements- Metals are solid, except for mercury, which is a liquid
- Most of the diatomic elements are gases: H2 (g),N2 (g), O2 (g), F2 (g), and Cl2 (g).
- Bromine is a liquid and iodine is a solid: Br2 (l) and I2 (s).
- Sulfur, phosphorus, and carbon are solids
Compounds
- All ionic compounds are solid at room temperature.
- An ionic compound that is very soluble is shown as aqueous when it is dissolved in water. An ionic compound that is slightly soluble is usually shown as a solid, even when it’s in water.
- Molecular compounds are very difficult to predict. The smaller the molecules are, the more they tend to be gases. The larger they are, the more they tend to be liquids and then solids.
Energy Changes
Exothermic ReactionsExothermic reaction – a chemical reaction that releases energy, usually in the form of heat, light, or electricity
Combustion – exothermic chemical reaction that occurs when oxygen reacts quickly with a substance to form a new substance or substances; burning
Endothermic Reactions
Endothermic reaction – a chemical reaction that is energy absorbing
Ex. barium hydroxide octahydrate + ammonium thiocyanate + energy → barium thiocyanate + ammonia + water
Biochemical Reactions
Photosynthesis – how plants convert sunlight to chemical energyie. energy + carbon dioxide + water → glucose + oxygen
Cellular respiration – how animals convert chemical energy into movement
ie. glucose + oxygen → energy + carbon dioxide + water
Characteristics of Chemical Reactions
All reactions involve the production of new substances with their own characteristic properties, including: state at room temperature, melting point, colour, and densityAll reactions involve the flow of energy, detected as a change in temperature during the reaction
When new substances form in chemical reactions sometimes changes of state can be observed: bubbles indicate a gas, precipitates indicate a solid
All chemical reactions are consistent with the law of conservation of mass
Conservation of Mass
Law of conservation of mass – total mass of the reactants in a chemical reaction equals the total mass of the productsAtoms cannot be created nor destroyed during a chemical reaction, they are simply rearranged into different molecules and sometimes different states.
Ex. When wood burns, the glucose in the wood reacts with the oxygen in air and produces carbon dioxide and water vapor, which are released into the air. If you were to take a large air tight container and burn the wood, the mass of the air and wood, before burning would be the same as the mass of the ashes, and smoke, after burning.
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A3.2 Writing Chemical Equations
Chemical equation – record of a chemical reaction using chemical symbols and formulas, shorthand way of showing the results of a chemical reaction
A chemical equation is very similar to a mathematical equation. The number of each type of atom must be the same on the left and right side of the arrow. It also includes the state of each compound.
Symbolizing Chemical Change
To write a chemical equation you require the following:
- Careful observation
- Knowledge of what substances are present at the start of the reaction
- The ability to analyze the materials produced by the reaction
Writing Word Equations
The reactants placed on the left side, and the products are placed on the right side of the arrow. A plus sign is placed between the names.
Writing Balanced Formula Equations
Formula equation – chemical equation that uses the chemical formulas of reactants and products in a chemical reaction
Skeleton equation – formula equation showing the identity of each substance involved in a chemical reaction; does not show the correct proportions of the reactants and the products
Keep the most complicated compound on the reactant side at one and balance all of the compounds
Balance the number of elemental molecules.
Ex. What is the balanced equation for the reaction between hydrogen and oxygen that forms water?
Word equation:
hydrogen + oxygen
→
water
Skeleton equation:
H2 + O2
→
H2O
Balanced equation:
2H2 + O2
→
2H2O
Try practice problem 1 on page 89
practice problem 1 on page 89
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A3.3 Five Common Types of Chemical Reactions
Formation Reactions
Formation reaction – chemical reaction in which two elements combine to form a compound; also known as a synthesis reaction
Synthesis reaction – chemical reaction in which two elements combine to form a compound; also known as a formation reaction
Ex.
carbon + oxygen → carbon dioxide
hydrogen + oxygen → water
Decomposition Reactions
Decomposition reaction – chemical reaction in which a compound breaks apart into its elements
Ex. When electricity runs through water the water molecule breaks into hydrogen gas and oxygen gas.
water → hydrogen + oxygen
Hydrocarbon Combustion
Hydrocarbons – compound that contains hydrogen and carbon; common hydrocarbons include the main components of gasoline and many plastics
Ex.
When methane, natural gas, is combusted, it forms carbon dioxide and water vapour.
methane + oxygen → carbon dioxide + water
Single Replacement Reactions
Single replacement reaction – chemical reaction in which a reactive element reacts with an ionic compound
Ex.
When sodium is added to water it produces sodium hydroxide and hydrogen.
sodium + water → sodium hydroxide + hydrogen
Double Replacement Reactions
Double replacement reaction – chemical reaction between two ionic compounds in solution that often results in the formation of a least one precipitate.
Ex.
When an acid is added to a base it produces a salt and water
hydrochloric acid + sodium hydroxide → sodium chloride + water
iron sulphide + hydrochloric acid → hydrogen sulphide + iron chloride
Predicting the Products of Chemical Reactions
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A3.4 The Mole
Avogadro’s Number and the Mole
Mole – quantity that chemists use to measure elements and compounds; symbol; mol; Avogadro’s number is the number of particles in a mole.
Avogadro’s number – number of atoms in 1 mole; approximately 6.02×1023; symbol NA
Molar Mass
Molar mass – mass of one mole of a substance
Atomic molar mass – average molar mass of an element’s atoms, including those of all the element’s different isotopes.
Chemists had to choose an element to be the standard for the conversion from atomic mass unit to grams. Since, carbon-12 is relatively easy to make 100% pure it was chosen. Carbon 12 has 6 protons and 6 neutrons so it has an atomic mass of 12 units. A mole of carbon 12 atoms has a mass of 12 grams.
The masses on the periodic table are weighted average mass for a mole of material. They are not exact, carbon is 12.01 because there are different forms of carbon, isotopes, so they took average samples.
Ex. How much does two moles of hydrogen chloride mass?
hydrogen chloride is HCl
H has a mass of 1.01 g/mol
Cl has a mass of 35.45 g/mol
HCl has a mass of 1.01 + 35.45 = 36.46 g/mol
2 HCl has a mass of 2 mol × 36.46 g/mol = 72.92 g
Ex. How many moles of water is in a 200g sample?
water is H2O
H has a mass of 1.01 g/mol
O has a mass of 16.00 g/mol
H2O has a mass of 18.02 g/mol
200 g ÷ 18.02 g/mol = 11.1 mole of water
Ex. What is the mass of 10 moles of sodium chloride?
sodium chloride is NaCl
Na has a mass of 22.99 g/mol
Cl has a mass of 35.45 g/mol
NaCl has a mass of 58.44 g/mol
10 moles of NaCl has a mass of 684.4 g
Ex. How many moles are in 2000 g of nitrogen gas?
nitrogen gas is N2
N has a mass of 14.01 g/mol
N2 has a mass of 28.02 g/mol
2000 g ÷ 28.02 g/mol = 111.0
The Mole Concept and the Law of Conservation of Mass
The mole concept and the law of conservation of mass allows chemists to predict the mass of substance produced.
Ex. the equation of the decomposition of water is
water → hydrogen + oxygen
2H2O → 2H2 + O2
2 moles of water will produce 2 moles of hydrogen gas and 1 mole of oxygen gas
36.04 g of water produces 4.04 g of hydrogen and 32 g of oxygen
Word equation: | hydrogen + oxygen | → | water |
Skeleton equation: | H2 + O2 | → | H2O |
Balanced equation: | 2H2 + O2 | → | 2H2O |
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Formation Reactions
Formation reaction – chemical reaction in which two elements combine to form a compound; also known as a synthesis reactionSynthesis reaction – chemical reaction in which two elements combine to form a compound; also known as a formation reaction
Ex.
carbon + oxygen → carbon dioxide
hydrogen + oxygen → water
hydrogen + oxygen → water
Decomposition Reactions
Decomposition reaction – chemical reaction in which a compound breaks apart into its elementsEx. When electricity runs through water the water molecule breaks into hydrogen gas and oxygen gas.
water → hydrogen + oxygen
Hydrocarbon Combustion
Hydrocarbons – compound that contains hydrogen and carbon; common hydrocarbons include the main components of gasoline and many plasticsEx.
When methane, natural gas, is combusted, it forms carbon dioxide and water vapour.
methane + oxygen → carbon dioxide + water
methane + oxygen → carbon dioxide + water
Single Replacement Reactions
Single replacement reaction – chemical reaction in which a reactive element reacts with an ionic compoundEx.
When sodium is added to water it produces sodium hydroxide and hydrogen.
sodium + water → sodium hydroxide + hydrogen
sodium + water → sodium hydroxide + hydrogen
Double Replacement Reactions
Double replacement reaction – chemical reaction between two ionic compounds in solution that often results in the formation of a least one precipitate.Ex.
When an acid is added to a base it produces a salt and water
hydrochloric acid + sodium hydroxide → sodium chloride + water
iron sulphide + hydrochloric acid → hydrogen sulphide + iron chloride
hydrochloric acid + sodium hydroxide → sodium chloride + water
iron sulphide + hydrochloric acid → hydrogen sulphide + iron chloride
Predicting the Products of Chemical Reactions
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A3.4 The Mole
Avogadro’s Number and the Mole
Mole – quantity that chemists use to measure elements and compounds; symbol; mol; Avogadro’s number is the number of particles in a mole.
Avogadro’s number – number of atoms in 1 mole; approximately 6.02×1023; symbol NA
Molar Mass
Molar mass – mass of one mole of a substance
Atomic molar mass – average molar mass of an element’s atoms, including those of all the element’s different isotopes.
Chemists had to choose an element to be the standard for the conversion from atomic mass unit to grams. Since, carbon-12 is relatively easy to make 100% pure it was chosen. Carbon 12 has 6 protons and 6 neutrons so it has an atomic mass of 12 units. A mole of carbon 12 atoms has a mass of 12 grams.
The masses on the periodic table are weighted average mass for a mole of material. They are not exact, carbon is 12.01 because there are different forms of carbon, isotopes, so they took average samples.
Ex. How much does two moles of hydrogen chloride mass?
hydrogen chloride is HCl
H has a mass of 1.01 g/mol
Cl has a mass of 35.45 g/mol
HCl has a mass of 1.01 + 35.45 = 36.46 g/mol
2 HCl has a mass of 2 mol × 36.46 g/mol = 72.92 g
Ex. How many moles of water is in a 200g sample?
water is H2O
H has a mass of 1.01 g/mol
O has a mass of 16.00 g/mol
H2O has a mass of 18.02 g/mol
200 g ÷ 18.02 g/mol = 11.1 mole of water
Ex. What is the mass of 10 moles of sodium chloride?
sodium chloride is NaCl
Na has a mass of 22.99 g/mol
Cl has a mass of 35.45 g/mol
NaCl has a mass of 58.44 g/mol
10 moles of NaCl has a mass of 684.4 g
Ex. How many moles are in 2000 g of nitrogen gas?
nitrogen gas is N2
N has a mass of 14.01 g/mol
N2 has a mass of 28.02 g/mol
2000 g ÷ 28.02 g/mol = 111.0
The Mole Concept and the Law of Conservation of Mass
The mole concept and the law of conservation of mass allows chemists to predict the mass of substance produced.
Ex. the equation of the decomposition of water is
water → hydrogen + oxygen
2H2O → 2H2 + O2
2 moles of water will produce 2 moles of hydrogen gas and 1 mole of oxygen gas
36.04 g of water produces 4.04 g of hydrogen and 32 g of oxygen
hydrogen chloride is HCl
H has a mass of 1.01 g/mol
Cl has a mass of 35.45 g/mol
HCl has a mass of 1.01 + 35.45 = 36.46 g/mol
2 HCl has a mass of 2 mol × 36.46 g/mol = 72.92 g
H has a mass of 1.01 g/mol
Cl has a mass of 35.45 g/mol
HCl has a mass of 1.01 + 35.45 = 36.46 g/mol
2 HCl has a mass of 2 mol × 36.46 g/mol = 72.92 g
water is H2O
H has a mass of 1.01 g/mol
O has a mass of 16.00 g/mol
H2O has a mass of 18.02 g/mol
200 g ÷ 18.02 g/mol = 11.1 mole of water
H has a mass of 1.01 g/mol
O has a mass of 16.00 g/mol
H2O has a mass of 18.02 g/mol
200 g ÷ 18.02 g/mol = 11.1 mole of water
sodium chloride is NaCl
Na has a mass of 22.99 g/mol
Cl has a mass of 35.45 g/mol
NaCl has a mass of 58.44 g/mol
10 moles of NaCl has a mass of 684.4 g
Na has a mass of 22.99 g/mol
Cl has a mass of 35.45 g/mol
NaCl has a mass of 58.44 g/mol
10 moles of NaCl has a mass of 684.4 g
nitrogen gas is N2
N has a mass of 14.01 g/mol
N2 has a mass of 28.02 g/mol
2000 g ÷ 28.02 g/mol = 111.0
N has a mass of 14.01 g/mol
N2 has a mass of 28.02 g/mol
2000 g ÷ 28.02 g/mol = 111.0
water → hydrogen + oxygen
2H2O → 2H2 + O2
2 moles of water will produce 2 moles of hydrogen gas and 1 mole of oxygen gas
36.04 g of water produces 4.04 g of hydrogen and 32 g of oxygen
2H2O → 2H2 + O2
2 moles of water will produce 2 moles of hydrogen gas and 1 mole of oxygen gas
36.04 g of water produces 4.04 g of hydrogen and 32 g of oxygen