Table of Contents

Chapter 1: States of Matter

Solids, Liquids & Gases

Solids:

Fixed volume, fixed shape, high density.

Atoms vibrate in place.

Closely packed together & regular pattern.

Liquids:

Fixed volume & take the shape of the container.

Less dense than solids, more dense than gas.

Particles move and slide past each other.

Gases:

Not fixed volume, take shape of the container.

Lowest density, a lot of space between particles.

So gases are easily compressed.

Random particles & particles move quickly.

Particles collide with each other & the sides of the container (this is how pressure is created).

State Changes

Melting

Solid → Liquid

Heat energy → Kinetic energy, and melting only occurs at a specific temp known as melting point.

Boiling

Liquid → Gas

Heat needed to form bubbles of gas below the surface of the liquid, allowing it to escape from the surface.

Occurs at boiling point (B.P.).

Freezing

Liquid → Solid

Reverse of melting, occurs at the same temperature.

Requires a lot of temperature decrease.

Evaporation

Liquid → Gas at any temperature.

Evaporation happens at the surface only.

High energy particles escape surface at any temperature.

The larger the surface area, the more quickly a liquid evaporates.

Condensation

Gas → Liquid at a range of temperatures.

When a gas is cooled, particles lose energy.

When they bump into each other, they don’t have enough energy to bounce away; instead, they group together and form a liquid.

Kinetic Theory

When a substance is heated, particles absorb thermal energy, which is converted to kinetic energy.

Heating solid causes particles to vibrate more, & as temperature increases, they vibrate to the point solid’s structure breaks & it melts.

With more heating, the liquid expands & rises until particles at the surface gain enough energy to escape & the liquid boils when they reach the gas phase.

These changes are shown on a heating curve & reverse is a cooling curve, these curves show how changes in temperature affect changes of state.

Pressure & Temperature in Gases

The more the temperature increases, the volume of gas also increases, and the density decreases if volume increases.

If you have gas stored in a container & squeezed, the pressure increases as you decrease the volume.

Gases & Kinetic Theory

Gaseous particles are in constant & random motion.

Pressure that gas creates inside a closed container is produced by gas hitting the walls of the container.

An increase in kinetic energy increases if temperature increases. With an increase in kinetic energy, particles move faster and collide with walls more frequently.

Decrease in volume = Increase in pressure.

Diffusion

Diffusion is the movement of particles from an area of high concentration to an area of low concentration, and eventually, the concentration of particles will be even as they spread out to occupy all the space.

Diffusion happens on its own & doesn’t require energy, but it happens faster at higher temperatures.

Diffusion occurs faster in gases compared to liquids. At the same temperature, gases don’t diffuse at the same rate.

This is due to the difference in molecular masses; lighter particles travel faster, therefore, they diffuse faster.

This can be demonstrated in the reaction between ammonia (NH₃) and hydrogen chloride gas (HCl) inside a glass tube. Where the two gases meet, white smoke of ammonium chloride (NH₄Cl) is formed.

However, they don’t meet in the middle of the tube; instead, they meet closer to the end where hydrogen chloride (Mr = 36.5) is. This is because ammonia (Mr = 17) molecules are lighter, so they travel faster than HCl.

Chapter 2: Atoms, Elements and Compounds

Elements, Compounds & Mixtures

Elements:

Substance made up of atoms that contain the same proton number & can’t be split into something simpler.

Compounds:

Substance made of two or more elements chemically combined. There’s an unlimited number of compounds. They can’t be separated.

Mixtures:

Mixture of two or more substances (elements and/or compounds) that aren’t chemically combined.

They can be separated by physical methods.

Atomic Structure & Periodic Table

All substances are made of tiny particles of matter called atoms.

An atom has subatomic particles – protons, neutrons, electrons.

The protons and neutrons are located at the center of the atom, which is called the nucleus.

Electrons move very fast around the nucleus in orbital paths called shells.

The mass of electrons is negligible; hence the mass of an atom is concentrated within the nucleus, where protons & neutrons are located.

Relative atomic mass unit is 1/12th the mass of a carbon-12 atom.

All elements are measured relative to the mass of a carbon-12 atom, so relative atomic mass has no units.

Relative Mass & Charge:

Proton: Mass 1, Charge positive 1+

Neutron: Mass 1, Charge no charge/neutral 0

Electron: Mass

$\frac{1}{1840}$ , Charge negative 1−

Proton Number:

Atomic number/proton number is the number of protons in the nucleus.

The symbol for atomic number is Z.

The number of electrons present in a neutral atom & determines the position of the element on the periodic table.

Mass Number:

Total number of protons & neutrons; the symbol is A.

Nucleon (mass) – proton number gives the neutron number.

Periodic table shows elements together with atomic number (proton number on top) & relative atomic mass at the bottom.

Note: Mr or Mass number isn’t the same.

In the periodic table:

Group VIII noble gases have a full outer shell = stable.

The number of outer shell electrons is equal to the group number of that element.

The number of occupied electron shells is equal to the period.

Electronic Configuration

Using diagrams called electron shell diagrams called the electronic configuration.

Electron Shell Diagrams

Electrons orbit the nucleus in shells (aka energy levels); further from the nucleus, there is more energy per shell.

Electrons fill the shell closest to the nucleus first.

First shell holds 2 electrons, 2nd shell holds 8 electrons, and for the first 20 elements, the third shell (above) can also hold 8, but this is a simplified view.

The outermost shell is called the valence shell.

The arrangement of electrons can also be explained by writing the number of electrons in each shell and separating by commas.

This is the electronic configuration.

Isotopes

Isotopes are different atoms of the same element that have the same number of protons but different numbers of neutrons.

Isotopes of the same element have the same chemical properties because they have the same number of electrons; therefore, the same electronic configuration.

The Isotopes Symbol:

It can also be written with chemical symbols and proton number & the mass number. E.g., C-14

Carbon – 6 electrons

Needs 4 electrons to stabilize, so when this carbon atom gains 4 electrons, the overall number of electrons increases over protons, so the charge is negative. e.g:

Relative Atomic Mass (RAM)

Relative Atomic Mass (RAM) is the average mass of all isotopes & mass number isn’t the same.

RAM takes isotopes into account while calculating.

Calculating Relative Atomic Mass

Symbol: Ar

The average mass of the isotopes of an element, compared to 1/12th of the mass of an atom of $^{12}\text{C}$

Relative atomic mass can be calculated from the mass number and relative abundances of all the isotopes using the equation:

$\text{Ar} = \frac{(\% \text{ of isotope 1 } \times \text{ mass number of isotope 1}) + (\% \text{ of isotope 2 } \times \text{ mass number of isotope 2})}{100}$

Top line can be extended to include lots of isotopes of an element.

Metallic Bonding

Metal atoms are held strongly by metallic bonding in a giant metallic lattice.

Within the metal lattice, the atoms lose electrons & become positively charged & the outer electrons no longer belong to a particular metal & are delocalized.

They move freely between the positive metal ions like a sea of electrons. Metallic bonds are strong & are a result of attraction between positive metal ions & negative charged delocalized electrons.

Properties of Metals:

High melting & boiling point, due to strong metallic bonds being hard to break.

Conduct electricity – have free electrons to move & carry charge.

Malleable & ductile – layers of positive ions can slide over each other & take different positions, strong but flexible so they can be made into specific shapes.

Ions & Ionic Bonds

An ion is an electrically charged atom or group of atoms formed by loss or gain of electrons. This is done by atoms to stabilize.

The loss/gain of electrons is done to gain full outer shell of electrons, which is a more stable arrangement of electrons.

Sodium Atom

There is space for 7 electrons, so it is easier to lose the existing electron instead of gaining 7.

Sodium Ion

There are more protons than electrons now that sodium lost an electron. Therefore, it is positively charged.

Metal:

All metals can lose electrons and become positively charged ions.

Metals, when losing electrons, have fewer electrons than protons, and hence become positively charged ions.

Non-Metal:

Non-metals can gain electrons and become negatively charged ions.

Non-metals, when gaining electrons, have more electrons than protons, and hence become negatively charged ions.

When metals react with non-metals, ionic compounds are made.

Metals lose their outer electrons, which the non-metals’ atoms gain.

Positive & negative ions are held together by strong electrostatic forces of attraction between opposite charges.

The force of attraction is known as the ionic bond.

Dot & Cross Diagrams:

Diagrams that show the arrangement of the outer-shell electrons in ionic or covalent compounds or elements.

e.g.:

Na (metal) & Cl (non – metal)

Lattice Structure of Ionic Compounds

Ionic compounds have a giant lattice structure.

Lattice structures are arranged in an ordered & repeating fashion.

The lattices of ionic compounds have a regular arrangement of alternating positive & negative ions.

Properties of Ionic Compounds

Usually solid at room temperature.

High melting & boiling points because of strong electrostatic forces; a lot of energy will be needed to break bonds.

Good conductors of electricity in molten or in solution; this is because the particles are charged & freely moving.

Poor conductors in solid state because they are in fixed positions within the lattice.

Simple Molecules and Covalent Bonds

Covalent compounds are formed when pairs of electrons are shared between atoms.

Only non-metal elements participate in covalent bonding.

When two or more atoms bond covalently, they are described as molecules.

Dot & cross diagrams can be used to show covalent bonds. E.g.,

Properties of Simple Molecular Compounds

Small compound molecules are covalently bonded. They have low melting & boiling points, so they are usually liquids or gases at room temperature.

As molecule size increases, so does melting & boiling point.

Small molecules are poor electricity conductors.

There is weak intermolecular force, so they have low melting & boiling points.

Intermolecular forces are weak, so most small molecules are liquid/gas.

As molecules increase in size, the intermolecular forces also increase; this causes melting & boiling point to increase.

Most molecular compounds are poor conductors as there are no free ions to carry charge, so most covalent compounds are insulators in solid state.

Giant Covalent Structures

Diamond & Graphite

Allotropes of carbon with giant covalent structures, due to differences in bond arrangement they are physically different in diamonds, carbon atoms bond with 4 other carbons forming a tetrahedron.

All bonds are identical, strong & have no intermolecular forces.

In graphite, carbon atoms bond to 3 others forming layers of hexagons, leaving one free electron per carbon to become delocalized.

The covalent bonds within layers are very strong, but the layers attract to each other with weak intermolecular forces.

Properties of:

Diamond

Insulator – no free electrons

High melting point & boiling point – strong bonds

Hard & dense

Used in jewellery & cutting tools.

Graphite

Has free delocalized electrons, so it can conduct electricity.

Layers are connected by weak forces, so they can slide over each other, so graphite is smooth & slippery.

High melting & boiling point.

Less dense than diamond.

Used in pencils & lubricants.

Used to make inert electrodes.

Silicon (IV) Oxide

A macromolecular compound which occurs naturally as sand & quartz.

Each oxygen atom covalently bonds with 2 silicon atoms & each silicon atom bonds with 4 oxygen atoms.

A tetrahedron is formed.

Comparing Diamond & Silicon

$SiO_2$ is hard & has a high melting point, insoluble in water & doesn’t conduct electricity.

$SiO_2$ is cheap as it’s naturally available & used in making sandpaper and to line the insides of furnaces.

Chapter 3: Stoichiometry

Formulae

Each element has its own unique symbol as in the periodic table.

Symbol contains two letters; the first is always capital & if there’s a second, it will be in lowercase.

Example: Sodium (Na)

Chemicals combine in fixed ratios that give full outer shells of electrons, chemical formulae tell the ratio of atoms, e.g.,

$H_{2} O$ – 2 hydrogen atoms combined with 1 oxygen.

Structural formula tells the way in which the atoms in a particular molecule are bonded. Done by a diagram or written version.

Molecular formula tells you the actual number of atoms of each element in one molecule of the compound or element.

Example: Butane – a compound

Structural Formula

Diagram:

$\text{H} – \text{C} – \text{C} – \text{C} – \text{C} – \text{H}$

Written:

$\text{CH}_3\text{CH}_2\text{CH}_2\text{CH}_3$

Molecular Formula

$\text{C}_4\text{H}_{10}$

Empirical Formula (Ratio)

$\text{C}_2\text{H}_5$

Simplest whole number ratio of atoms of each element in a molecule.

Organic molecules have different empirical & molecular formulae.

Formula of ionic compound is always an empirical formula.

Valency

The number of electrons to lose/gain for a full outer shell. You can use valency to find chemical formula:

Aluminium sulfide

Al valency = 3 S valency = 2

$\text{Al}_2\text{S}_3$

Deducing Formulae of Ionic Compounds

The valency can tell you the charge, based on loss/gain. e.g.: H needs to lose, so it is $1^-$ charge, and bromide needs to gain an electron, so the charge is $1^-$ .

The overall sum of charges must be 0.

Example: Sodium Bromide

Na Br

$1^+$ $1^-$

1 sodium needed for one bromide.

$\text{NaBr}$

Writing Word Equations & Symbol Equations

Word Equations:

Show reactants & products using full chemical names.

Arrow implies conversion of reactants to products.

Reaction conditions or catalysts can be written above the arrow.

Example:

Sodium hydroxide + hydrochloric acid → Sodium chloride + water.

Names of Compounds

For compounds with metal & non-metal, the metal comes first & the second atom has -ide at the end.

If both are non-metals, and one is hydrogen, hydrogen comes first.

For other non-metals, the element with the lower group number comes first.

For compounds with atoms that occur regularly, e.g., Carbonate $(\text{CO}_3^{2-})$ , Hydroxide $(\text{OH}^-)$ , Sulfate $(\text{SO}_4^{2-})$ , and Nitrate $(\text{NO}_3^-)$ . When these ions form compound with a metal atom, the metal name comes first.

Examples:

$KOH$

$(Potassium Hydroxide)$

$\text{CaCO}_3$

(Calcium Carbonate)

Writing and Balancing Chemical Equations

Same number of atoms on each side needed.

$H$ , $N$ , $F$ , $O$ , $Cl$ , $Br$ , and I must be diatomic, e.g.,

$\text{H}_2 + \text{O}_2$ Groups of atoms like Nitrate $({NO}_{3}^{-})$ that hasn’t changed from one side to another, count it as a group, not individual atoms.

State Symbols

Written after each formula to show the physical state of each substance.

Aqueous means a substance is dissolved in water.

Examples:

Solid (s)

Liquid (l)

Gas (g)

Aqueous (aq)

Deducing Symbol Equations

Step 1: Work out formula and state symbols of reactants and products to construct an unbalanced symbol equation.

Step 2: Balance the equation.

Balancing Ionic Equations

In aqueous solutions, ionic compounds *dissociate into their ions.

*Separate into the component ions that formed them.

Steps:

Write balanced equation.

Identify ionic substances & write down the ions separately.

Rewrite by eliminating ions that appear on both sides of the equation.

Example:

$2 K I (a q) + C l_{2} (a q) \to 2 K C l (a q) + I_{2} (a q)$

$2K^+(aq) + 2I^-(aq) + Cl_2(aq) \rightarrow 2K^+(aq) + 2Cl^-(aq) + I_2(aq)$

$2I^-(aq) + Cl_2(aq) \rightarrow 2Cl^-(aq) + I_2(aq)$

Relative Masses of Atoms & Molecules

Relative Atomic Mass, Ar:

The average mass of isotopes of an element compared to 1/12th of the mass of a $^{12}\text{C}$

atom.

Relative Molecular Mass, Mr:

The sum of relative atomic masses; they are used for ionic compounds.

Calculating Reacting Masses

Mass can’t be created or destroyed.

Example:

$2\text{Ca} + \text{O}_2 \rightarrow 2\text{CaO}$

Relative atomic mass of Ca = 40, O = 16.

This means 80 units $(2 \times 40)$ $of Ca react with$ $(2 \times 16)$ 32 units of oxygen.

$2 Ca + O_{2} \to 2 CaO$

80 + 32 = 112.

The ratio of mass will always be the same regardless of the unit, e.g., 80g Ca reacts with 32g of O will produce 112g of CaO. Or 80 tonnes of Ca with 32 tonnes of oxygen will produce 112 tonnes of calcium oxide.

The Mole and Avogadro’s Constant

The mole, mol, is the unit of amount of substance that contains $6.02 \times 10^{23}$ particles.

The number $6.02 \times 10^{23}$ is known as Avogadro’s Constant.

Amount of Substance (mol):

$\text{mol} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

The mass of 1 mole of substance is known as molar mass – it is the same as Ar/Mr.

$\text{Mols} = \frac{\text{Grams (g)}}{\text{RAMs (relative atomic mass)}}$

Reacting Masses & Limiting Reactants

Chemical equations can be used to calculate the moles/mass of reactants & products.

Information in question is used to find the amount in moles, then the ratio is identified through a balanced chemical equation.

Once moles are determined, they can be converted into grams using Ar or Mr.

The reaction stops when a reactant is used up, the reactant used up first is the limiting reactant; it limits the duration & the amount of product produced.

The product is directly proportional to the limiting reactant.

The limiting reactant is the one not present in excess.

To Perform Reacting Mass Calculations:

Write balanced equation.

Calculate the moles of each reactant.

Compare the moles & deduce the limiting reactant.

Concentration

Concentration simply refers to the amount of solute there is in a specific volume of solvent. The unit for concentration is g/dm³ or mol/dm³.

Concentration (g/dm³):

$Concentration = \frac{mass of solute (g)}{volume of solution (dm³)}$

$\text{Concentration} = \frac{\text{mass of solute (g)}}{\text{volume of solution (dm³)}}$ $1\text{ dm}^3 = 1 \text{ litre}$

A common constant to use is the molar gas volume which is $24 {dm}^{3}$ at room temperature & pressure (R.T.P.) for gas calculations.

Titration Calculations

Titration is a method of analyzing the concentration of solutions. Acid-base titrations are one of the most important titration kinds. You can be asked to calculate moles present in a given amount, the concentration, or the volume required to neutralize acid/base.

Example:

Step 1: Write equation for reaction:

$HCl (aq) + NaOH (aq) \to NaCl (aq) + H_{2} O (l)$

Step 2: Calculate the moles for NaOH (Substance neutralizing the hydrochloric acid).

Step 3: Deduce the moles of acid – use ratio, concentration & moles of NaOH.

Step 4: Find concentration of acid – use the mols & volume (given in Q).

Empirical Formula

The simplest whole number ratio of the atoms of each element present in one molecule or formula unit.

Organic molecules often have different empirical & molecular formulae. The formula of an ionic compound is always an empirical formula.

Molecular Formula

The actual number of atoms of each element present.

To calculate molecular formula:

Step 1: Find the relative formula mass of the empirical formula.

Step 2:

$\frac{\text{Relative formula mass of molecular formula}}{\text{Relative formula mass of empirical formula}}$

Step 3: Multiply the number of each element present in the empirical formula by the number in step 2.

Percentage Yield

Yield describes the amount of product you get from a reaction.

You can never get 100% yield due to several reasons, e.g., reactant left behind in equipment, product lost during transfer, etc.

Actual yield is the recorded amount of product obtained.

Theoretical yield is the amount of product that would be obtained under perfect conditions.

$% Yield = \frac{Actual Yield}{Theoretical Yield} \times 100$

Percentage Composition by Mass

$% of element = \frac{Total mass of element in compound}{Relative formula mass of compound} \times 100$

Percentage Purity

$\% \text{Purity} = \frac{\text{Mass of pure substance}}{\text{Total mass of substance}} \times 100$

Chapter 4: Electrochemistry

Electrolysis

Electrolysis is the decomposition of an ionic compound, when molten or in aqueous solution, by the passage of electric current.

The transfer of charge during electrolysis includes:

Movement of electrons in the external circuit.

Loss or gain of electrons at the electrodes.

The movement of ions in the electrolyte.

During electrolysis, metals or hydrogen are formed at the cathode, and non-metals are formed at the anode.

Metal objects are electroplated to improve their appearance and resistance to corrosion.

Metals are electroplated by:

Place an object at the cathode – the object to be electroplated.
Place pure metal at the anode – the metal will coat the object at the cathode.
The electrolyte is an aqueous solution of soluble salt of pure metal at the anode.

At the anode, metal loses electrons & forms ions in solution, at the cathode, ions gain electrons & become tin atoms that deposit on the object, thus electroplating it.

Electroplating is used to make metals resistant to corrosion & to improve the appearance.

Identifying the Products

Binary Compound in Molten State

Binary compound is a compound with 2 different elements joined ionically. To predict their products made at electrodes, identify the ions:

Positive ion to the cathode, negative to the anode. Cathode always has metal product & anode has non-metal. Excluding Hydrogen!!

Examples:

Lead (II) Bromide

Pb $^{2+}$ br $^{-}$

anode – brown gas – bromide

Cathode – grey metal – lead

Potassium Chloride

K $^{+}$ Cl $^{-}$

Anode – green/yellow gas – chlorine

Cathode – metal build up – potassium

Halide Compound in Concentrated/Aqueous Solution

Aqueous NaCl

Brine is a concentrated solution of aqueous sodium chloride.

When electrolyzed, gas is formed at both electrodes (chlorine & hydrogen).

Both are important:

Chlorine for bleach.

Hydrogen for margarine (butter-like spread).

Sodium hydroxide for soap & detergents.

At Cathode

H discharged (H $^{+}$ gains electron).

Hydrogen is more reactive than sodium.

At Anode

Cl gas discharged.

Cl loses electrons, Na $^{+}$ & OH $^-$ remain in solution & form NaOH.

Dilute H $_2$ SO $_4$

At Cathode:

H $^+$ at cathode, hydrogen gas released.

At Anode:

OH $^-$ at anode, oxygen gas & water released.

To Identify Gas Released:

Oxygen – glowing splint test.

Chlorine – bleaches damp litmus paper to white.

Hydrogen – pop squeak test.

Aqueous Copper (II) Sulfate with Inert & Copper Electrodes

Inert (graphite/carbon):

Oxygen at anode.

Carbon electrodes dissolve, so the oxygen becomes CO $_2$ .

Cu $^{2+}$ at the cathode.

Copper:

Copper deposits on the cathode.

Anode dissolves, copper is oxidized to ions & dissolves into the solution.

Hydrogen-Oxygen Fuel Cells

A hydrogen-oxygen fuel cell uses hydrogen and oxygen to produce electricity with water as the only chemical product.

Advantages & Disadvantages of Using Hydrogen-Oxygen Fuel Cells Rather than Gasoline/Petrol Vehicles:

Advantages	Disadvantages
Don’t produce pollution.	Fuel cells are expensive.
Release more energy.	Difficult & expensive to store.
No power lost in transmission.	Fuel cells are affected by low temperatures, less efficient.
Quieter, less noise pollution.	Fewer hydrogen filling stations.

Ionic Half-Equations

Oxidation: Substance loses electrons.

Reduction: Substance gains electrons.

As the ions in the electrolyte come in contact with the electrodes, electrons aren’t lost or gained, and they form neutral substances, which are discharged as products at the end.

At the anode: Negative ions lose electrons & are oxidized.

At the cathode: Positive ions gain electrons & are reduced.

Ionic half-equations show oxidation & reduction of ions & result in balanced charges.

Examples:

Molten Lead Bromide

At anode: $2Br^-$ → $Br_2$ + $2e^-$

At cathode: $Pb^{2+}$ + $2e^-$ → $Pb$

Dilute Sulfuric Acid

At anode: $4OH^-$ → $O_2$ + $2H_2O$ + $4e^-$

At cathode: $2H^+$ + $2e^-$ → $H_2$

OIL: Oxidation is loss of electrons.

RIG: Reduction is gain of electrons.

Chapter 5: Chemical Energetics

Exothermic Reaction

Exothermic reactions transfer thermal energy to the surroundings, leading to an increase in temperature of the surroundings.

Endothermic Reaction

Endothermic reactions take in thermal energy from the surroundings, leading to a decrease in temperature of the surroundings.

Enthalpy Change

For atoms to react with each other in a chemical system, they must come in contact through a collision.

Number of factors affect collisions: energy, orientation, number of collisions per second (frequency of collisions).

There is a minimum amount of energy required for a collision to be successful; that is, for particles to react.

Transfer of Energy

The transfer of thermal energy during a reaction is called the enthalpy change (ΔH) of the reaction. ΔH is negative for exothermic reactions and ΔH is positive for endothermic reactions.

Activation energy, $E_a$ , is the minimum energy that colliding particles must have to react (successful collision).

Pathway Diagrams

Bonds & Enthalpy

Bond breaking is an endothermic process, and bond making is an exothermic reaction.

The Enthalpy Change of a Reaction

Each chemical bond has specific bond energy associated with it; this is the amount of energy required to break the bond or the amount of energy given out when the bond is formed.

This energy can be used to calculate how much heat is released or absorbed during the reaction.

Method of Calculating Bonds

Write a balanced equation if none are present.

Add all the bond energies for the bonds in the reactants – energy in.

Add all the bond energies for the bonds in the products – energy out.

Calculate the enthalpy change.

Enthalpy Change (ΔH) = Energy In – Energy Out.

Chapter 6: Chemical Reactions

Physical & Chemical Changes

Physical Changes:

Easy to reverse; observable.

Examples: melting, evaporating, etc.

Chemical Changes:

New chemical substance formed; very difficult to reverse.

Exothermic:

Give out energy/heat.

Endothermic:

Take in energy/heat.

Rate of Reaction

Collision Theory

In order for a reaction to happen, particles must collide with each other and have sufficient energy to cause a successful collision – reaction.

The minimum energy required for a successful collision is called Activation Energy (Ea).

The number of successful collisions depends on the number of particles per unit volume, the frequency of collisions, and kinetic energy of particles.

The particles must have sufficient activation energy.

Effect on the Rate of Reaction of:

Concentration of Solution – Increased = higher rate of reaction.

Higher concentration = more particles = more collisions = faster reaction.

Pressure of Gases – Increased = higher rate of reaction.

Higher pressure = more collisions = faster reaction.

Surface Area of Solid – Increased = higher rate of reaction.

Increased = more places available for the reaction to occur faster.

Temperature – Increased = higher rate of reaction.

Increased temperature = more kinetic energy = more collisions = faster reaction.

Adding/Removing Catalyst – Adding = speeds up reaction.

Adding catalyst decreases $E_a$ needed so more successful collisions lead to a faster reaction.

A catalyst increases the rate of reaction of a reaction without being used up or changed. A catalyst decreases the activation energy $E_{a}$ of a reaction.

Experimenting

Method Choice	Advantage	Disadvantage
Disappearing cross experiment	Simple experiment with no specialist equipment required	Difficult to determine when the cross is obscured. Different people will determine the cross to have disappeared at different levels of cloudiness.
Collecting gas using a gas syringe	All the gas collected is from the reaction. Gas syringes are easy to set up and keep gas dry.	Gas syringes are expensive and fragile. The pistons can stick. Limited volume can be collected. Gas can be lost if the bung is not put onto the top of the flask quickly.
Collecting gas using an inverted measuring cylinder	The experiment uses common lab equipment.	The delivery tube beneath the measuring cylinder can come away. It can be difficult to read the scale, especially when the reaction is vigorous with a large number of bubbles produced rapidly. Gas can be lost if the bung is not put onto the top of the flask quickly.
Measuring the loss of mass of reactants	Simple experiment to carry out with no specialist equipment	Not suitable for gases with a small Mr.

Evaluating/Interpreting the Graphs

The steeper the curve, the faster the rate of the reaction.

The curve is steepest initially, so the rate is quickest at the beginning of the reaction.

As the reaction progresses, the concentration of the reactants decreases, and the rate decreases, shown by the curve becoming less steep.

When one of the reactants is used up, the reaction stops, the rate becomes zero, and the curve levels off to a horizontal line.

The amount of product formed in a reaction is determined by the limiting reactant:

If the amount of limiting reactant increases, the amount of product formed increases.

If the amount of the reactant in excess increases, the amount of product remains the same.

Drawing a tangent to the slope allows you to show the gradient at any point on the curve.

The steeper the slope, the quicker the rate of reaction.

Redox

Roman numerals are used to indicate oxidation number (the number of electrons gained/lost during forming a chemical bond) of an element in a compound, e.g., iron(II) oxide (iron loses 2 electrons).

A redox reaction is a reaction that involves simultaneous oxidation and reduction.

Oxidation	Reduction
Gain of oxygen	Loss of oxygen
Loss of electrons	Gain of electrons
Increase in oxidation number	Decrease in oxidation number

Example:

Zinc oxide + carbon → zinc + carbon monoxide

$ZnO + C \to Zn + CO$

Zinc oxide has been reduced, it lost oxygen.

Carbon has been oxidized, it gained oxygen.

Zinc + copper sulfate → zinc sulfate + copper

$Zn + {CuSO}_{4} \to {ZnSO}_{4} + Cu$

Spectator ion: remove $\text{SO}_4^{2-}$ because it doesn’t change.

$Zn(s) + {Cu}^{2 +} (aq) \to {Zn}^{2 +} (aq) + Cu(s)$

Write half-equations:

$\text{Zn (s)} \rightarrow \text{Zn}^{2+} (\text{aq}) + 2e^-$ (oxidized, lost electrons).

$\text{Cu}^{2+} (\text{aq}) + 2e^- \rightarrow \text{Cu (s)}$ (reduced, gained electrons).

The oxidation number of elements in their uncombined state is zero.

The oxidation number of a monatomic ion is the same as the charge on the ion.

The sum of oxidation numbers in a compound is zero.

The sum of oxidation numbers in an ion is equal to the charge of the ion.

Identifying Redox Reactions by Color Changes

Acidified Potassium Manganate ${KMnO}_{4}$ $\text{KMnO}_4$	Potassium Iodide $KI$ $\text{KI}$
Oxidizing agent	Reducing agent
Used to test for presence of a reducing agent.	Used to test for oxidizing agents.
When acidified potassium manganate is added to a reducing agent, the color changes from purple to colorless.	When potassium iodide is added to an oxidising agent, the color changes from colorless to red – brown color.

Oxidising Agent: A substance that oxidizes another & reduces itself.

Reducing Agent: A substance that reduces another & oxidizes itself.

Example:

Copper (II) oxide + Hydrogen → Copper + Water

$CuO + H_{2} \to Cu + H_{2} O$

${Cu}^{2 +} + 2 e^{-} \to Cu gain electrons,$ $reduced oxidising$ $agent$

$H_{2} \to 2 H^{+} + 2 e^{-} loses electrons,$ $oxidized, reducing$ $agent$

Reversible & Equilibrium

Some chemical reactions are reversible and are shown with the symbol $⇌.$

A reversible reaction in a closed system is at equilibrium when the rate of the forward reaction is equal to the rate of the reverse reaction, and the concentration of reactants and products are no longer changing.

The position of equilibrium is affected by changing temperature because increasing temperature will shift equilibrium in the endothermic direction. Increased pressure shifts equilibrium in the direction that produces fewer molecules of gas. Increased concentration shifts equilibrium to the right to reduce the effect of the increase. A catalyst does not affect the position of equilibrium but it increases the rate at which equilibrium is reached.

Copper (II) Sulfate

${CuSO}_{4}$ $\text{CuSO}_4$	${CuSO}_{4}$ $\text{CuSO}_4$
It loses water of crystallization as a result of evaporation, and the crystals turn from blue to white and from crystalline form to amorphous (no clear shape/form).	Anhydrous copper sulfate is a white powder that turns blue in contact with water.
	$\text{CuSO}_4 \cdot 5\text{H}_2\text{O} \rightleftharpoons \text{CuSO}_4 + 5\text{H}_2\text{O}$

Cobalt (II) Chloride $\text{CuSO}_4$

The Effect of Heat on Hydrated ${CoCl}_{2}$	Addition of Water to Anhydrous ${CoCl}_{2}$
When heated, the pink crystals turn to blue anhydrous cobalt (II) chloride.	Anhydrous $\text{CoCl}_2$ is blue, and when water is added, it turns pink.
${CoCl}_{2} \cdot 6 H_{2} O ⇌ {CoCl}_{2} (s) + 6 H_{2} O(l)$

The Haber Process

$\text{N}_2 (g) + 3\text{H}_2 (g) \rightleftharpoons 2\text{NH}_3 (g)$

The Haber process is used to produce ammonia for industrial use. Ammonia is used for fertilizers, manufacturing plastic, etc.

The nitrogen in the Haber process is from the air, and hydrogen from methane gas.

The typical conditions for the Haber process are:

450°C – Forward is exothermic, increasing temperature shifts the position to the left, lowering yield.

20,000 kPa / 200 atm to shift equilibrium to the right; too high pressure is dangerous.

Iron beads are used as a catalyst to help reach equilibrium faster.

The Stages of the Haber Process:

Step 1: $\text{H}_2$ & $\text{N}_2$ are obtained through natural gas and pumped into a compressor through a pipe.

Step 2: The gases are compressed to 200 atm inside the compressor.

Step 3: The pressurized gases are pumped into a tank with iron beads at a temperature of 450°C.

Step 4: Some $\text{H}_2$ & $\text{N}_2$ react, the unreacted gas $\text{NH}_3$ passes to a cooling tank. $\text{NH}_3$ is liquefied and removed to storage vessels.