1. States of matter

Definitions:

1. Solid: In solids, particles are held in fixed positions. This is because of the strong forces of attraction between particles in solids. Particles in a solid vibrate around their fixed positions

2. Liquid: In liquids, particles are randomly arranged and free to move past each other. The particles never stop moving. This is because of the weak forces of attraction between particles in liquids

3. Gas: In gases, particles are and free to move and far away from each other. This is because of the very weak forces of attraction between particles in gases. Particles are constantly moving.

Changes of state:

• Solids turn to liquids when they melt.

  • The increase in kinetic energy causes particles to vibrate more.

  • At a certain temperature, particles break free from their positions.

  • Liquids turn back into solids during freezing.

• Liquids turn to gases when they evaporate.

  • Further increases in kinetic energy cause particles to move faster.

  • At a certain temperature, particles break their bonds.

  • Gases turn back into liquids during condensation.

• Solids can turn directly into gases during sublimation.

  • If enough kinetic energy is supplied, substances can bypass the liquid phase.

Solutions and solubility

Definitions:

• Diffusion: is the movement of particles from an area of high concentration to an area of lower concentration

• Dilution: is lowering the concentration of a solute in a solution by simply adding more solvent to the solution

• A solvent: is a liquid in which a solute dissolves

• A solute: is the substance being dissolved - it is usually a solid

• A solution: is the mixture formed when a solute dissolves in a solvent

• Saturated: when no more solute can dissolve into a solution

• Solubility: how many grams of a solute can dissolve in 100g of water

Things that affect solubility:

• Temperature increases → solubility increases (solids)

• Pressure increases → solubility increases (gasses)

• The solvent

2. Element, compounds and mixtures

Definitions + properties:

1. Elements: consist of only one type of atom.

   • For example, hydrogen (H2), nitrogen (N2) and carbon (C) are all elements.

2. Compounds: are made up of two or more different elements.

   • These elements are chemically bonded to one another.

   • For example carbon (C) and oxygen (O2) can react with each other to form carbon dioxide (CO2).

   • The properties of compounds are often very different to their original elements

3. Mixtures: are a combination of two or more different elements.

   • Mixtures are not chemically bonded to one another.

   • Instead they are physically in close proximity.

   • For example, sand and water can form a mixture.

   • Mixtures will show the properties of both elements

Pure substances: are made out of only one element or compound

• Pure substances have melting points and boiling points that will not change

Methods of Separation:

1. Distillation: is used to separate out solutions

   1. Simple distillation: only works when the components have very different boiling points

      1. 1. The solution is heated.

         • The part of the solution with the lowest boiling point will evaporate.

      2. The vapor is cooled and condensed.

         • This is done with the aid of a cooling water jacket.

      3. The newly condensed, pure liquid is collected in a new container

      4. The rest of the solution remains in the flask.

   

   1. Fractional distillation: used to separate mixtures of liquids

      1. These liquids will have different boiling points.

      2. The liquid with the lowest boiling point will evaporate first.

      3. The evaporating liquid will rise up a fractionating column.

      4. It will then be carried to a condenser.

         • Here it will reform as a liquid.

      5. The temperature is then raised so the second liquid can be collected.

2. Filtration: used to separate insoluble solid from a liquid

   1. The filter paper is placed in a funnel above a beaker.

   2. The mixture is poured into the funnel.

3. Crystallization: separates soluble solids from a solution

   1. The solution is heated in an evaporating dish.

      • Some of the liquid evaporates, creating a more concentrated solution

   2. As the solution cools, crystals of solute begin to form

      • This is because the solute is much less soluble when it is cold.

   3. The crystals are then filtrated from the solution and dried

4. Paper chromatography: is a way to separate mixtures

   • The mixtures are separated by their solubility in that particular solvent

   1. A pencil line is drawn close to the bottom of the filter paper

      • The pencil is insoluble so it won’t dissolve in the solvent

   2. Spots of different inks are placed equidistant from each other on the pencil line

   3. The paper is lowered into the solvent but it should not touch the inks (below the pencil line the solvent should be)

   4. As the solvent goes up the filter paper, it carries the dye with it

3. Atomic structure

Definitions:

• An atom is the smallest particle of an element.

  • An atom contains three subatomic particles:

    • Protons - these are heavy and positively charged.

    • Neutrons - these are heavy and neutral.

    • Electrons - these have hardly any mass and are negatively charged

Structure of an atom:

• The protons and the neutrons are found in the nucleus

• The electrons are found on the other shells of the atom

Atoms V.S. Molecules

Terms:

• Atomic number: number of protons

• Mass number: number of protons + number of neutrons

• Isotopes: two (or more) atoms of the same type that have different mass numbers (same number of protons but different neutrons number)

• Relative atomic mass: The average mass of naturally occurring atoms of an element on a scale where the carbon 12 atom has a mass of exactly 12 units. Basically since atoms have such a small mass, the atoms mass is compared to the ratio between carbon 12 where its relative atomic mass is 12

How to calculate the relative atomic mass of an element:

4. The periodic table

• Elements are arranged in order of atomic (proton) number

Groups: These are the vertical columns that show how many outer electrons each atom has and are numbered from 1 – 7, with a final group called group 0 (instead of group 8)

Periods: These are the horizontal rows that show the number of shells of electrons an atom has and are numbered from 1 - 7

What is the electron configuration of an atom?

The electronic configuration of an element tells you how many electrons are in each shell around an electron’s nucleus

e.g. sodium has 11 electrons: 2 in its most inner shell, then 8, then 1 in its outermost shell

Metals and non-metals in the periodic table:

Metals = elements that react to form positive ions.

• Majority of elements are metals.

  • Found to the left and towards the bottom of the periodic table.

• Non-metals = elements that do not form positive ions.

  • Found towards the right and top of the periodic table

Why doesn’t group 0 react?

• They are unreactive and do not easily form molecules, because they have a stable arrangement of electrons (always have full outer shells)

5. Chemical formulae

1. Calculating relative mass

What is Mr (Relative Molecular Mass) in chemistry?

• We have seen previously that the symbol for the relative atomic mass is Ar

• This is calculated from the mass number and relative abundances of all the isotopes of a particular element

• To calculate the Mr of a substance, you have to add up the relative atomic masses of all the atoms present in the formula

Calculating percentage by mass of an element in a compound:

Example:

2. Calculating Moles, Mass & RFM

The Mole:

• Chemical amounts are measured in moles

• The symbol for the unit mole is mol

• One mole of a substance contains the same number of the stated particles, atoms, molecules, or ions as one mole of any other substance

Molar mass: One mole of any element is equal to the relative atomic mass of that element in grams or for a compound the relative formula mass in grams

Example: How many moles are in 2.64 g of sucrose, C12H22O11 (Mr = 342.3)?

Answer:

1. The number of moles is found by mass ÷ molar mass

2. This comes to 2.64 g ÷ 342.3 g mol-1 = 7.71 x 10-3 mol

3. Calculating reactive masses

• The information given in the question is used to find the amount in moles of the substances being considered

• Then, the ratio between the substances is identified using the balanced chemical equation

• Once the moles have been determined they can then be converted into grams using the relative atomic or relative formula masses

Example:

Calculate the mass of magnesium oxide that can be made by completely burning 6.0 g of magnesium in oxygen in the following reaction:

Balancing Equations using Reacting Masses

• If the masses of reactants and products of a reaction are known then we can use them to write a balanced equation for that reaction

• This is done by converting the masses to moles and simplifying to find the molar ratios

Example: A student reacts 1.2 g of carbon with 16.2 g of zinc oxide. The resulting products are 4.4 g of carbon dioxide and 13 g of zinc. Determine the balanced equation for the reaction.

4. Calculating percentage yield

Yield is the term used to describe the amount of product you get from a reaction

Example: Copper(II) sulfate may be prepared by the reaction of dilute sulfuric acid with copper(II) oxide. A student prepared 1.6 g of dry copper(II) sulfate crystals.

Calculate the percentage yield if the theoretical yield is 2.0 g.

• Actual yield of copper(II) sulfate = 1.6 g

• Percentage yield of copper(II) sulfate = 1.6/2.0 × 100 = 80%

5. Empirical & Molecular Formulae

Empirical formula: gives the simplest whole number ratio of atoms of each element in the compound

e.g.

6. Calculate Concentrations of Solutions

Example: Calculate the amount of solute, in moles, present in 2.5 dm³ of a solution whose concentration is 0.2 mol dm^-3.

7. Calculate Volumes of Gases

• At room temperature and pressure, the volume occupied by one mole of any gas was found to be 24 dm³ or 24,000 cm³

  • This is known as the molar gas volume at RTP (room temperature and pressure)

6. Ionic bonding

Definition: Ions are atoms that have lost or gained electrons

• Metal atoms lose electrons to become positively charged ions

• Nonmetal atoms gain electrons to become negatively charged ions

Terms:

• Cation: positive ion

• Anion: negative ion

Definition: Ionic compounds form when metals give one or more electrons to non-metals.

The strong attraction between a cation and an anion is called an electrostatic attraction. The forces act in all directions in the lattice.

Properties of ionic bonding:

• Ionic compounds have high melting and boiling points.

• Ionic compounds are hard and brittle.

• Ionic compounds dissociate into ions when dissolved in water.

• Solutions of ionic compounds and melted ionic compounds conduct electricity, but solid materials do not.

Reason for these properties:

• There are strong electrostatic forces of attraction between oppositely charged ions

• Requires a lot of energy to overcome these forces of attraction

7. Covalent bonding

Definition: Covalent bonding occurs in most non-metallic elements and in compounds of nonmetals

How the bonding works:

Strong bonds between atoms that are covalently bonded are the result of electrostatic attraction between the positive nuclei of the atoms and the pairs of negative electrons that are shared between them.

1. Simple molecular structures

   • They are are attracted to each other through intermolecular forces

   • Substances with a simple molecular structures are gases or liquids, or solids with low melting and boiling points

   • This is because even if the covalent bonds are strong, between the compounds the intermolecular forces are weak

2. Giant covalent structures

   • Substances that consist of giant covalent structures are solids with very high

     melting and boiling points.

   • All of the atoms in these structures are linked to other atoms by strong covalent

     bonds

   • Examples:

     • Diamond

       • very hard

       • doesn’t conduct electricity

     • Graphite

       • soft

       • conducts electricity

Properties:

• do not conduct electricity (except graphite)

8. Metallic bonding

Metals consist of giant structures of atoms arranged in a regular pattern

• The electrons in the outer shell of metal atoms are delocalized and so are free to move through the whole structure.

• The sharing of delocalized electrons makes a strong electrostatic attraction between negatively charged electrons and positive metal ions

Properties:

• Most metals have high melting and boiling points.

• They can conduct heat and electricity because of the delocalized electrons in their structures.

• The layers of atoms in metals are able to slide over each other, so metals can be bent and shaped

9. Electrolysis

• When an electric current is passed through a molten ionic compound the compound decomposes or breaks down

• The process also occurs for aqueous solutions of ionic compounds

• Covalent compounds cannot conduct electricity hence they do not undergo electrolysis

• Ionic compounds in the solid state cannot conduct electricity either since they have no free ions that can move and carry the charge

Key terms:

• Electrode: is a rod of metal or graphite through which an electric current flows into or out of an

  electrolyte

• Electrolyte: is the ionic compound in molten or dissolved solution that conducts the electricity

• Anode: is the positive electrode of an electrolysis cell

• Anion: is a negatively charged ion which is attracted to the anode

• Cathode: is the negative electrode of an electrolysis cell

• Cation: is a positively charged ion which is attracted to the cathode

Positive electrode (anode)

• Non-metal ions (other than hydrogen) are attracted to the positive electrode

• Non-metal ions will lose electrons to form the non-metal

• The product formed depends on which ion loses electrons more readily, with the more reactive ion remaining in solution

Negative electrode (cathode)

• Hydrogen and metal ions attracted to the negative electrode but only one will gain electrons

• Either hydrogen or metal will be produced

• If the metal is above hydrogen in reactivity series, then hydrogen will be produced and bubbling will be seen at the cathode

Electrolysis of Aqueous Solutions

• Aqueous solutions will always have water present (H2O)

• Which ions get discharged and at which electrode depends on the relative reactivity of the

  elements involved

  • Anode:

    • If halide ions (Cl- , Br- , I- ) and OH- are present then the halide ion is discharged at the anode, loses electrons and forms a halogen (chlorine, bromine or iodine)

    • If no halide ions are present, then OH is discharged at the anode, loses electrons and forms

      oxygen gas

  • Cathode:

    • If the metal is above hydrogen in the reactivity series, then hydrogen will be produced and bubbling will be seen at the cathode

    • This is because the more reactive ions will remain in the solution, causing the least reactive ion to be discharged

Gas tests:

• Hydrogen: squeaky pop test

  • The test for hydrogen consists of holding a burning splint held at the open end of a test tube of gas.

  • If the gas is hydrogen it burns with a loud “squeaky pop” which is the result of the rapid combustion of hydrogen with oxygen to produce water.

• Chlorine: Litmus paper

  • Chlorine is an acidic gas that also acts as a bleach. Damp

    litmus paper is bleached white when it is placed in chlorine.

  • If damp blue litmus paper is used, the paper turns red then white.

• Carbon dioxide: Limewater

  • Carbon dioxide reacts with calcium hydroxide solution to produce a white precipitate of calcium carbonate.

  • Limewater is a solution of calcium hydroxide. If carbon dioxide is bubbled through limewater, the limewater turns milky or cloudy white.

• Oxygen: glowing splint relights

  • Oxygen supports combustion. If oxygen is present in a test tube, a glowing splint relights when it is held inside.

Reactions at the Electrodes & Ionic Half-Equations:

• Oxidation is when a substance loses electrons

• Reduction is when a substance gains electrons

• At the anode, negatively charged ions lose electrons and are thus oxidized

• At the cathode, the positively charged ions gain electrons and are thus reduced

Half equations show the oxidation and reduction of the ions involved

Transfer of Charge

1. During electrolysis the electrons move from the power supply towards the cathode

2. Positive ions within the electrolyte move towards the negatively charged electrode which is the cathode

3. Here they accept electrons from the cathode and either a metal or hydrogen gas is produced

4. Negative ions within the electrolyte move towards the positively charged electrode which is the anode

5. If the anode is inert (such as graphite or platinum), the ions lose electrons to the anode and form a nonmetal or oxygen gas

6. If the anode is a reactive metal, then the metal atoms of the anode lose electrons and go into solution as ions, thinning the anode

Electroplating: is a process where the surface of one metal is coated with a layer of a different metal

• The anode is made from the pure metal you want to coat your object with

• The cathode is the object to be electroplated

Uses of electroplating

1. Electroplating is done to make metals more resistant to corrosion or damage

   • e.g, chromium and nickel plating

2. It is also done to improve the appearance of metals,

   • e.g. coating cutlery and jewelry with silver

Conductors:

• Solids such as metals or graphite

• Liquids such as molten lead bromide or molten metals

• Solutions such as sodium chloride solution

Insulators: Most insulators are solids of plastic, rubber or ceramic