The Three States of Matter

Properties of the States of Matter

Properties of Solids

• Regular arrangement

• They DONT move, they vibrate.

• They are not fluids and cannot be compressed.

• Particles are very dense(closely packed).

• Low energy between particles

Properties of Liquids

• Irregular arrangement

• The particles slide past each other

• Can be compressed

• Particles are close

• Medium energy between particles

• Fluid

Properties of Gases

• Random arrangement

• Move quickly and randomly in all directions

• Can be compressed

• Fluid

• Particles are far

• High energy between particles

Phase Changes

Melting

• Melting is when a solid changes into a liquid

• The process requires heat energy which transforms into kinetic energy, allowing the particles to move

• It occurs at a specific temperature known as the melting point which is unique to each pure solid

Boiling

• Boiling is when a liquid changes into a gas

• This requires heat which causes bubbles of gas to form below the surface of a liquid, allowing for liquid particles to escape from the surface and from within the liquid

• It occurs at a specific temperature known as the boiling point which is unique to each pure liquid

Freezing

• Freezing is when a liquid changes into a solid

• This is the reverse of melting and occurs at exactly the same temperature as melting, hence the melting point and freezing point of a pure substance are the same

  • Water for example freezes and melts at 0 ºC

• It requires a significant decrease in temperature (or loss of thermal energy) and occurs at a specific temperature which is unique for each pure substance

Evaporation

• When a liquid changes into a gas

• Evaporation occurs only at the surface of liquids where high energy particles can escape from the liquids surface at low temperatures, below the boiling point of the liquid

• The larger the surface area and the warmer the liquid/surface, the more quickly a liquid can evaporate

• Evaporation occurs over a range of temperatures, but heating will speed up the process as particles need energy to escape from the surface

Condensation

• When a gas changes into a liquid, usually on cooling

• When a gas is cooled its particles lose energy and when they bump into each other, they lack energy to bounce away again, instead grouping together to form a liquid

Sublimation

• When a solid changes directly into a gas

• This happens to only a few solids, such as iodine or solid carbon dioxide ( dry ice )

• The reverse reaction also happens and is called desublimation or deposition

Diffusion and Dilution

Definitions

Diffusion

Diffusion is the spreading out of particles in a gas or liquid. There is an over all movement of particles from areas of high concentration to low concentration.

The bromine gas (Br2) becomes more pale as it diffuses into the air due to less concentration of particles.

Examples of diffusion:

Spraying perfume/deoderant

Farting

Aqueous - Dissolved in water. Written as (aq) in chemical equations.

Gaseous - Written as (g) in chemical equations.

Solid - Written as (s) in chemical equations.

Liquid - Written as (l) in chemical equations

Ammonia - NH3 - Colourless

NH3(g) + HCl(g) = NH4Cl(s)

These compounds form into the solid ring ring NH4Cl via diffusion.

The ring of ammonium chloride forms nearer to the HCl because the particles of ammonia are smaller and lighter and diffuse through the air more quickly.

Gas particles dont instantly form a solid because the may collide with air particles, and they move in random directions.

Dilution

If you pour water into orange juice, the orange colour fades. This is dilution. The more you dilute a liquid, the more the colour fades away. This is because it increases the concentration of the solute in the solvent.

It is decreasing concentration of the solute by increasing the volume of the solvent.

Solutions Terminology

Water is a universal solvent.

Solubility

Solubility is a measurement of how much of a substance will dissolve in a given volume of liquid.

The solubility of the solvent is dependent on pressure and temperature.

Different substances have different solubilities

Solubility can be expressed in g per 100 g of solvent

Solubility of solids is affected by temperature

As temperature increases, solids usually become more soluble

Solubility of gases is affected by temperature and pressure; in general:

As pressure increases, gases become more soluble

As temperature increases, gases become less soluble

Solubility = (Mass of solute dissolved / Mass of water)* 100

Solubility graphs or curves represent solubility in g per 100 g of water plotted against temperature

To plot a solubility curve, the maximum mass of solvent that can be dissolved in 100 g of water before a saturated solution is formed, is determined at a series of different temperatures

Elements, Compounds and Mixtures

Definitions

Element - Made up of only one type of atom. The simplest type of substances. - Oxygen

Compounds - A substance that contains two or more different elements chemically bonded in fixed proportions. - Water

Molecule - Two chemically bonded atoms that can be the same or different. Chemically bonded in fixed proportions

Mixture - Different substances in the same space but not chemically bonded. - Salt Water

Diatomic Molecule - An element that needs two of the same atom to be chemically stable - Oxygen

HNOF Down: Diatomic Molecules

(Go from Nitrogen to Fluorine on the periodic table then down; all of these elements are diatomic. And hydrogen!)

Pure Substances vs Mixture

Pure substances melt/boil at a fixed temperature while mixtures melt/boil over a range of temperatures.

Foreign substances in a crystalline solid disrupt the repeating pattern of forces that holds the solid together. Therefore, a smaller amount of energy is required to melt the part of the solid surrounding the impurity

Pure Substance

Mixture (Impure Substance)

Separation Techniques

Simple Distillation

Simple distillation is used to separate soluble solute inside of mixtures while keeping the solvent.

1. Heat the flask until the solvent in the solution evaporates.

2. The solvent will travel through the condenser and condense into it’s pure form, while the solute is left behind.

Water enters through the bottom of the condenser. This is because if water entered at the top of the condenser, then the water would always be leaving the condenser; never keeping it full.

Simple distillation is used when the solute and solvent have different boiling points.

Fractional Distillation

Fractional distillation is used to separate two miscible liquids.

Miscible - Solution where a liquid completely dissolves in another liquid.

1. The solution is heated to the temperature of the substance with the lowest boiling point.

2. This substance will rise and evaporate first, and vapours will pass through a condenser, where they cool and condense, turning into a liquid that will be collected in a beaker.

3. All of the substance is evaporated and collected, leaving behind the other components(s) of the mixture.

Filtration

Filtration is used when an insoluble solid is to be separated from a mixture of the solid and the solvent.

A piece of filter paper is placed in a filter funnel above a beaker

A mixture of insoluble solid and liquid is poured into the filter funnel

The filter paper will only allow small liquid particles to pass through as filtrate

Solid particles are too large to pass through the filter paper so will stay behind as a residue

Crystallisation

Crystallisation is used to separate a dissolved solute from a solvent while getting rid of the solvent. If you wanted to keep both the solvent and solute, you would use simple distillation.

The solution is heated, allowing the solvent to evaporate, leaving a saturated solution behind

Test if the solution is saturated by dipping a clean, dry, cold glass rod into the solution

If the solution is saturated, crystals will form on the glass rod

The saturated solution is allowed to cool slowly

Crystals begin to grow as solids will come out of solution due to decreasing solubility

The crystals are collected by filtering the solution, they are washed with cold distilled water to remove impurities and are then allowed to dry

Paper Chromatography

Rf - Retention factor

This technique is used to separate a mixtures / substances that have different solubilities in a given solvent.

A pencil line is drawn on chromatography paper and spots of the sample are placed on it. Pencil is used for this as ink is soluble.

The paper is then lowered into the solvent container, making sure that the pencil line sits above the level of the solvent, so the samples don’t wash into the solvent container

The paper is called the stationary phase

The solvent travels up the paper by capillary action, taking some of the coloured substances with it; this is called the mobile phase

Different substances have different solubilities so will travel at different rates, causing the substances to spread apart

The substances with higher solubility will travel further than the others

This will show the different components of the mixture / substance

If two or more substances are the same, they will produce identical chromatograms

If the substance is a mixture, it will separate on the paper to show all the different components as separate spots

An impure substance will show up with more than one spot, a pure substance should only show up with one spot

Rf = Distance travelled by solute / Distance travelled by solvent

Atoms: Definitions and Structure

Atom - The simplest type of substance. Building blocks of the universe.

Molecule - Two chemically bonded atoms that can be the same or different.

One relative atomic mass = 1/12 of a carbon12 atom.

Atoms are made up of sub-atomic particles; electrons, neutrons and protons.

Number of Neutrons = Atomic mass - Atomic number

Relative Atomic Mass - The ratio of the average mass per atom of an element to 1/12 of the mass of an atom of carbon-12.

Calculate Relative Atomic Mass

Periodic Table

Elements are arranged in increasing atomic / proton number

Groups are the amount of valence electrons each element as

Periods are the amount of electron shells there are

Electronic Configurations

Calcium, an element with 20 electrons has a configuration of 2,8,8,2

Amount of electrons in each shell

Classify Metals & Non-Metals

Electronic Configuration & Reactivity

• Elements in the same group in the periodic table will have similar chemical properties

• This is because they have the same number of outer electrons so will react and bond similarly

• The group number of an element which is given on the periodic table indicates the number of electrons in the outer shell (valence electrons)

  • This rule holds true for all elements except helium; although is in group 0, it has only one shell, the first and innermost shell, which holds only 2 electrons

• We can use the group number to predict how elements will react as the number of valence shell electrons in an element influences how the element reacts.

• Therefore, elements in the same group react similarly

  • By observing the reaction of one element from a group, you can predict how the other elements in that group will react

  • By reacting two or more elements from the same group and observing what happens in those reactions you can make predictions about reactivity and establish trends in reactivity in that group

• For example, lithium, sodium and potassium are in group 1 and can all react with elements in group 7 to form an ionic compound

• The group 1 metals become more reactive as you move down the group while the group 7 elements show a decrease in reactivity moving down the group

• The elements in group 0 of the periodic table are called the noble gases

• They are all non-metal, monatomic (exist as single atoms), colourless, non-flammable gases at room temperature

• The group 0 elements all have full outer shells of electrons; this electronic configuration is extremely stable

• Elements participate in reactions to complete their outer shells by losing, gaining, or sharing electrons

  • The Group 0 elements do not need to do this, because of their full outer shells which makes them unreactive and inert

• Other than helium which has 2 electrons in its outer shell, the noble gases have eight valence electrons (which is why you may see this group labelled “group 8”)

  • Electronic configurations of the Noble gases:

    • He = 2

    • Ne = 2, 8

    • Ar = 2, 8, 8

    • Kr = 2, 8, 18, 8

    • Xe = 2, 8, 18, 18, 8

Ionic Bonding

Ion - An atom with a positive or negative charge. The atom loses or gains an electron to have a full outermost shell.

Anion - Negatively charged ion

Cation - Positively charged ion

All metals lose electrons and form cations. ( Except transition metals )

All non-metals gain electrons and form anions.

Diagram showing the formation of the sodium ion

Diagram showing the formation of the chloride ion

Common Ions

Ag+= Silver

NH4+= Ammonium

H+ = Hydrogen

OH- = Hydroxide

NO3- = Nitrate

Fe2+ = Iron(II)

Pb2+ = Lead

Zn2+ = Zinc

Cu2+ = Copper

SO4 2- = Sulfate

CO3 2- = Carbonate

Fe 3+ = Iron(III)

PO4 3- = Phosphate

Facts

Ionic compounds are normally solid at room temperature and are non-volatile (doesn’t vaporise easily).

They are usually water soluble because ionic compounds and water are polar substances.

Ionic compounds can conduct electricity when molten. The giant lattice structure is broken and ions are free to move around and carry charge.

Covalent Bonding

Covalent Bonding - The electrostatic forces of attractions between the oppositely charged positive nuclei and shared pair of electrons.

Simple Molecular Structure - Covalent bonds joining the non-metal atoms together but intermolecular bonds that act between neighboring molecules. This makes the structure weaker than a giant covalent structure.

Giant covalent structure - Covalent bonds joining all non-metal atoms together. High boiling and melting point

Viscosity - How thick a liquid is

Broken intermolecular forces are a physical change

Broken covalent bonds are a chemical change

The higher the relative molecular mass, the higher the boiling / melting point of a molecule

Covalent structures do not conduct electricity.

Only graphite does.

Diamond, graphite and sand are giant covalent structures. Diamond is composed of only carbon and each carbon atom has 4 covalent bonds. Sand is silicon dioxide. Each silicon atom is bonded to three oxygen atoms.

Graphite is made of layers of carbon atoms. To connect these layers there are intermolecular forces.

Each carbon atom has 3 covalent bonds, leaving one delocalised electron. It is able to move freely and carry charge, and is why graphite can conduct electricity.

Graphite is soft as the layers can slide over each other.

Metallic Bonding

Metallic Bonding - The electrostatic forces between cations and delocalised electrons.

Metals are good conductors of heat and electricity. This is because the delocalised electrons move freely and can store electrical currents and thermal energy

Alloys are mixtures of metals, where the metals are mixed together physically but are not chemically combined

They can also be made from metals mixed with nonmetals such as carbon

Alloys often have properties that can be very different to the metals they contain, for example they can have greater strength, hardness or resistance to corrosion or extreme temperatures

Alloys contain atoms of different sizes, which distorts the regular arrangements of atoms

This makes it more difficult for the layers to slide over each other, so they are usually much harder than the pure metal