Cambridge IGCSE Combined and Co-ordinated Sciences - C2.01 The states of matter

C2.01 The states of matter

• Matter:

  • Definition: Anything that has mass and occupies space (volume).

  • Chemistry studies matter's behavior and transformations.

• States of Matter:

  • Three states: solid, liquid, and gas.

  • State changes occur with temperature and/or pressure variations.

• General Characteristics (Table C2.01):

  • Solid:

    • Fixed volume and definite shape.

    • High density.

    • Does not flow.

  • Liquid:

    • Fixed volume but no definite shape (takes the shape of the container).

    • Moderate to high density.

    • Generally flows easily.

  • Gas:

    • No fixed volume (expands to fill the container).

    • No definite shape (takes the shape of the container).

    • Low density.

    • Flows easily.

  • Liquids and gases are fluids.

• Response to Temperature and Pressure:

  • All states expand with increased temperature and contract with decreased temperature.

  • Gases exhibit the most significant volume change with temperature.

  • Gases are easily compressible; liquids are slightly compressible; solids are nearly incompressible.

Changes in Physical State

• State Changes (Figure C2.01):

  • Melting: Solid to liquid.

  • Freezing (Solidification): Liquid to solid.

  • Evaporation: Liquid to gas.

  • Condensation (Liquefaction): Gas to liquid.

  • Sublimation: Solid to gas.

  • Reverse Sublimation (Deposition): Gas to solid.

• Melting and Freezing:

  • Melting Point (m.p.): Temperature at which a pure substance changes from solid to liquid.

    • Occurs at a specific temperature for each pure substance.

  • Freezing Point: Temperature at which a liquid turns into a solid.

    • The freezing point is the same as the melting point for a given substance.

    • Example: Water's melting/freezing point is $0°C$.

  • Gallium: A metal with a melting point slightly above room temperature; it melts in hand.

• Sublimation:

  • Direct conversion from solid to gas (e.g., carbon dioxide - dry ice).

    • Occurs at a specific temperature for each pure solid.

    • Example: Iodine sublimes into a purple vapor and condenses on cold surfaces.

• Evaporation, Boiling, and Condensation:

  • Evaporation: Liquid changes to gas at the surface.

    • Occurs over a range of temperatures.

    • Rate increases with surface area and temperature.

  • Boiling: Gas forms within the liquid at a specific temperature.

    • Boiling Point (b.p.): Specific temperature for each pure liquid (Figure C2.02).

  • Volatile Liquid: Evaporates easily and has a relatively low boiling point.

    • Example: Ethanol (b.p. $78 °C$) is more volatile than water (b.p. $100 °C$).
      *Condensation: The reverse of evaporation, usually caused by cooling or increasing pressure.

Pure Substances

• Definition: Consists of only one substance without impurities.

  • Melts and boils at definite temperatures.

• Melting and Boiling Points Table (C2.02):

  • Lists melting and boiling points of common substances.

  • Used to determine the state of a substance at room temperature (approximately $20 °C$).

• Purity Testing:

  • Melting and boiling points are precise and predictable for pure substances.

  • Used to test the purity of a sample or check the identity of an unknown substance.

• Melting Point Measurement (Figure C2.03):

  • Uses a melting-point apparatus with a capillary tube containing a powdered solid.

  • A water bath (for temperatures below $100 °C$) or an oil bath (above $100 °C$) heats the tube.

• Temperature Note:

  • Remember that $-100 °C$ is a higher temperature than $-150 °C$.

The Effect of Impurities and Purification

• Impure Substances:

  • Example: Seawater contains salt residue after evaporation.

  • Freezes below $0 °C$ and boils above $100 °C$.

  • Melting or boiling occurs over a range of temperatures.

• Effects of Impurities:

  • Lowers the melting point.

  • Raises the boiling point.

• Heating and Cooling Curves:

  • Heating Curve (Figure C2.04): Shows temperature change during heating.

    • Pure substances (e.g., naphthalene) melt sharply at a specific temperature.

    • Mixtures (e.g., wax) melt over a range of temperatures.

  • Cooling Curve (Figure C2.05): Shows temperature change during cooling.

    • Temperature remains constant during condensation and freezing.

    • Cooling mixtures (ice and salt) can lower the temperature below $0 °C$.

  • Energy Changes: Heat energy is needed to change a solid into a liquid or a liquid into a gas, and heat energy is given out during the reverse processes.

Types of Mixture

• Mixtures:

  • Made from at least two parts (solid, liquid, or gas).

  • States can be completely mixed (solution) or remain separate (suspension).

• Solutions:

  • Solid dissolves in liquid (e.g., salt in water).

  • Solute: The solid that dissolves.

  • Solvent: The liquid in which the solid dissolves.

  • Solute particles are dispersed completely and cannot be seen.

• Other Solutions:

  • Dissolved gases in seawater (oxygen and carbon dioxide).

  • Liquids in liquids (alcohol and water are miscible).

  • Alloys: Mixtures of metals (dissolving one metal in another before solidification).

C2.02 Separating and purifying substances

• Separation Techniques:

  • Essential for understanding and utilizing substances in mixtures.

• Factors for Choosing a Separation Method:

  • Type of mixture.

  • The substance of interest.

• Separation Methods (Table C2.03):

  • Solid + Solid (powdered mixture): Use differences in properties (density, solubility, sublimation, magnetism).

  • Suspension of Solid in Liquid: Filtration or centrifugation.

  • Liquid + Liquid (immiscible): Separating funnel or decantation.

  • Solution of Solid in Liquid:

    • To obtain solid: Evaporation (crystallization).

    • To obtain liquid: Distillation.

  • Two (or more) Liquids Mixed Together (miscible): Fractional distillation.

  • Two (or more) Solids in a Liquid: Chromatography.

• Separating Insoluble Solids from Liquids:

  • Decanting: Pouring off the liquid after the solid settles.

  • Filtration (Figure C2.07a):

    • Residue: Solid collected on filter paper.

    • Filtrate: Liquid collected.

  • Buchner Funnel (Figure C2.07b):

    • Uses a vacuum pump to speed up filtration.

  • Centrifugation: Spinning the mixture at high speed to deposit the solid at the bottom.

• Separating Immiscible Liquids:

  • Using a separating funnel, liquids separate into layers.

  • The denser layer is tapped off at the bottom.

• Distillation (Figure C2.08):

  • Used to separate a liquid from a solution.

  • The liquid evaporates and is then condensed by cooling.

  • Simple Distillation: Used when the boiling points of the substances are very different.

• Fractional Distillation (Figure C2.10):

  • Used to separate miscible liquids with different boiling points.

  • Fractionating column: Glass beads provide a large surface area for repeated evaporation and condensation.

  • Process: The mixture is heated, and vapors pass up the fractionating column. The substance with the lower boiling point evaporates, is cooled, and collected.

Separating two or more dissolved solids

*Chromatography: Seperates tow or more dissolved solids in a solution
*Chromatography: Can tell us whether a solution has become contaminated.
*Chromatography types: They all follow the same basic principles.

*Paper Chromatography:
*Drop of concentraded solution is placed on a pencil line near the bottom edge of a strip of chromatography paper.
*The level of the solvent must start below the sample.
*Solvent begins to move up the paper by capillary action.
*The solvent moves up the other components along at different rates.
*The separation of the mixture is complete.
*The different components string out along the paper like runners in a race.

Different Solvents Use in Chromatography:

*Water and organic solvents are common
*Organic solvents are useful because they dissolve many substances that are insoluble in water.

*The substances separate according to their solubility in the solvent.
*The substance that is most soluble moves fastest ip the paper.
*An insoluble substance would remain at the origin.

*The distance moved by a particular spot is measured and is related to the position of the solvent front.

*The ratio of these distances is called the R, value, or retention factor.
$R_f= \frac{distance \,moved \,by \,the \,substance}{distance \,moved \,by \,the \,solvent \,front}$

Locating Agents and Biological Molecules Analysis

*Locating agents allow method to be used for seperating substances that are not colored.
*Locating Agent: Paper is treated with to react with the samples to produce colored spots
*Chromatography: Very uselful in the analysis of biologically important molecules such as sugars, amino acids and nucleotide bases

*UV Light: Amino acids can be 'seen' if the paper chromatogram is viewed under ultraviolet light.

*Purity: If the sample is pure, it should only give one spot when run in several different solvents

General Purity Tests and Pharmaceuticals

*Measurments Of Melting Point or Boiling Point: Most gernally used tests for purity
*Impurities: lower the melting point or raise the boiling point of the substance.

*Pharmaceuticals: Must be of highest possible purity. Any contaminating substance, even in very small ammounts, may have harmful side effects.
*Food Colourings: Need to be carefully controlled (E100 to E180).

*Tartrazine (E102): May cause hyperactivity and allergic reactions to some children, for example asthma

Solubility

*A solution is made up of:
*The solute: The solid that dissolves.
*The solvent: The liquid in which it dissolves.
*Water is the commonest solvent in use, but other liquids are also important.
*If a substance dissolves in a solvent, it is said to soluble: if it does not dissolve, it is insoluble.
*Concentrated solution: Contains a high proportion of solute
*Dilute Solution: Contains a small proportion of solute.
*Saturated Solution: A solution that when more when more solid is added, no more will dissolve.

*The concentration of a solution is the mass of solute dissolved in a particular volume of solvent, usually $1dm^3$

*The solubility increases with temperature.

*Crystallisation: When a saturated solution is cooled it can hold less solute at the lower temperature, and some solute crystallises out.

Solubility of Gases in Liquids

*Solubility: Unlike most solids, gases become less soluble in water as the temperature rises.
*Oxygen: Is more soluble in water than nitrogen is.
*61% nitrogen and 37% oxygen when air is dissolved in water.
Gases: Solubility increases with pressure.
*Sparkling Drinks: Contain carbon dioxide dissolved under pressure
*Carbon Dioxide: Is more soluble than either nitrogen or oxygen.
*Carbonic Acid: Carbon dioxide reacts with water

C2.03 Atoms and molecules

*PURE SUBSTANCES:
*Elements: Substances that cannot be chemically broken down into simpler substances
*Compounds: pure substances made from two, or more elements chemically combined together

*Elements:
*Are the 'building blocks' from which the Universe is constructed.
*Over 100 known elements, but most of the Universe consists of just two.
*Hydrogen (92%) and helium (7%) make up most of the mass of the Universe, with all the other elements contributing only 1% to the total
*94 elements found naturally on Earth eight account for more than 98% of the mass of the Earth's crust.
*Two elements, silicon and oxygen, which are bound together in silicate rocks, make up almost three-quarters of the crust.
*The human body contains 65% oxygen, 18% carbon, 10% hydrogen, 3% nitrogen, 2% calcium and 2% of other elements.

*Chemical Reactions and Physical Changes:
*In a chemical reaction, a substance can be transformed (changed) into another substance.
*Decomposition is a chemical reaction, where a compound breaks down to form two or more substances
*Syntesis is the opposite where the substance is formed by the combination of two or more other substances

Comparing Mixtures and Compounds

*Atomic Theory: Based on the atomic theory put forward by John Dalton in 1807

*Dalton's theory:
*a pure element is composed of atoms
*the atoms of each element are different in size and mass
*atoms are the smallest particles that take part in a chemical reaction
*atoms of different elements can combine to make molecules of a compound.

*The kinetic model:
*All matter is made up of very small particles
*Articles are moving all the time
*The freedom of movement and the arrangement of the Particles is different for the three states of matter
*sure of a gas is produced by the atoms Couples of the gas hitting the walls of the container

*Key Points of Kinetic Theory:
*Lighter particles move more quickly than heavier particles at the same temperature; larger molecules diffuse more slowly than smaller ones.
*The pressure of a gas is the result of collisions of the fast-moving particles with the walls of the container.
*The average speed of the particles increases with an increase in temperature.
*H₂O
*CH
*HCI

2.04 The structure of the atom

Atoms: Are made up of sub-atomic particles-protons, neutrons and electrons.

Sub-Atomic Particles

• Protons:

  • Relative mass: 1

  • Relative charge: +1

  • Location: Nucleus

• Neutrons:

  • Relative mass: 1

  • Relative charge: 0

  • Location: Nucleus

• Electrons:

  • Relative mass: $\frac{1}{1840}$, negligible

  • Relative charge: -1

  • Location: Outside nucleus

Atomic Structure

• Electrically neutral atom: Equal numbers of protons and electrons.

• Hydrogen (H): 1 proton, 0 neutrons, 1 electron

• Helium (He): 2 protons, 2 neutrons, 2 electrons

Atomic and Mass numbers

*Atomic number: Number of protons in the nucleus
$Z = number \, of \, protons \, in \, the \, nucleus$
*Mass (Nucleon) number: Number of protons and neutrons in the nucleus
$A = number \, of \, protons + number \, of \, neutrons$

Element Notation

• Element symbol with atomic and mass numbers: $^A<em>ZX$ (e.g., $^7</em>3Li$)

Useful relationships

$Number \, of \, electrons = Number \, of \, protons = Atomic \, number \, (Z)$
$Number \, of \, Neutrons = Nucleon \, Number - \, Proton \, Number = A-Z$

TABLE 2;08 SHOWS RULE APPLES EUVEN TO MOST COMPLICATED ATOM

Remember that you can use the periodic table you have in exam for information on these numbers for an atom

• Different atoms of an element with the same number of protons but different numbers of neutrons.

  • Different masses due to varying neutron numbers.

Electron Configuration

• Niels Bohr's Theory (1913):

  • Electrons orbit the nucleus in shells (energy levels).

  • Shells further from the nucleus have higher energy.

  • Shells fill starting with the lowest energy level (closest to the nucleus).

  • First shell: Maximum 2 electrons.

  • Second and subsequent shells: Maximum 8 electrons for stability (noble gas configuration).