Chemistry for CSEC - States of Matter and Atomic Structure

Chemistry

Introduction to Chemistry

• Chemistry: The study of the structure and behavior of matter.

• Matter: Anything that has mass and occupies space.

• The three main states of matter are solid, liquid, and gas.

A1.1 The Particulate Nature of Matter

• Democritus's idea: Matter consists of particles.

• Particulate Theory of Matter: All matter is made of particles.

• This theory explains the physical properties of matter and the differences between the three states of matter.

• Four Main Ideas of the Particulate Theory:

  • All matter is made of particles.

  • The particles are in constant, random motion.

  • There are spaces between the particles.

  • There are forces of attraction between the particles.

• Examples explained by the particulate theory:

  • Density differences between solids, liquids, and gases.

  • Cooling a liquid to form a solid.

  • The movement of smells in a room.

  • Pressure increase in a gas with increased temperature.

  • Crisper vegetables when soaked in water.

  • Surface tension in liquids.

A1.2 Evidence for the Particulate Theory of Matter

• Diffusion: Movement of particles from an area of higher concentration to an area of lower concentration until evenly distributed.

• Smells traveling through the air are an example of diffusion.

• Practical Activity: Investigating diffusion in gases

  • Ammonia gas and hydrogen chloride gas diffuse towards each other in a glass tube.

  • They react to form a white solid, ammonium chloride (NH_4Cl(s)).

  • The reaction occurs closer to the hydrochloric acid because ammonia particles are lighter and move faster.

  • NH3(g) + HCl(g) \rightarrow NH4Cl(s)

• Osmosis: Movement of water molecules from a region with a lot of water molecules to a region with few water molecules through a differentially permeable membrane.

• A differentially permeable membrane allows some substances to pass through but not others.

• Practical Activity: Investigating osmosis in green paw-paw.

  • Cell membranes act as differentially permeable membranes.

  • Water moves into cells in distilled water, making the paw-paw more rigid and longer.

  • Water moves out of cells in concentrated sodium chloride, making the paw-paw floppy, softer, and shorter.

Practical Uses of Osmosis

• Controlling Garden Pests: Sodium chloride is sprinkled on slugs and snails. It absorbs moisture from their bodies, causing dehydration and death.

• Preserving Food Items: Sodium chloride and sugar withdraw water from cells of food items (meat, fish, fruits), preventing decay by osmosis. This also prevents the growth of decay-causing microorganisms.

A1.3 The Three States of Matter

• Matter exists in three states: solid, liquid, and gas.

• Physical Properties: Characteristics observed or measured without changing the chemical composition (e.g., shape, volume, density).

• The states differ in energy and arrangement of particles.

  • Solids: Particles vibrate in fixed positions, packed closely together, least energy.

  • Liquids: Particles move slowly, small spaces between them, medium energy.

  • Gases: Particles move rapidly, large spaces between them, greatest energy.

• The energy of particles relates to temperature; temperature changes can cause changes of state (physical change).

• Changes of State: Melting, evaporation, boiling, condensation, freezing, and sublimation occur with temperature changes.

• Melting: Solid becomes liquid when heated; particles gain kinetic energy. The temperature remains constant at the melting point.

• Evaporation: Liquid particles near the surface gain enough kinetic energy to become a vapor, cooling the liquid.

• Boiling: Liquid particles gain enough kinetic energy to change into gas both within the liquid and at the surface at the boiling point.

  • Boiling occurs at a specific temperature whereas evaporation can take place at any temperature.

  • Boiling takes place throughout the liquid, whereas evaporation takes place only at the surface of the liquid.

• Condensation: Gas becomes liquid when temperature lowers; particles lose kinetic energy.

• Freezing: Liquid becomes solid when temperature lowers; particles lose kinetic energy. The temperature at which a liquid freezes is called freezing point.

• Sublimation: Solid changes directly into gas (or vice versa) without passing through the liquid state.

• Examples of substances which undergo sublimation are iodine, carbon dioxide (known as 'dry ice'), ammonium chloride and naphthalene.

• Heating Curves: Temperature of a substance is measured at intervals as heat is added and changes state from solid to liquid to gas.

• Cooling Curves: Temperature of a substance is measured at intervals as heat is removed and changes state from gas to liquid to solid.

Mixtures and Their Separation

A2.1 Elements, Compounds, and Mixtures

• Pure Substances: Fixed composition, fixed properties, cannot be separated by physical means.

• Mixtures: Variable composition, components retain individual properties, can be separated by physical means.

• Elements: Cannot be broken down by ordinary chemical or physical means. Smallest particle is an atom.

• Compounds: Two or more elements chemically bonded in fixed proportions, cannot be separated by physical means. Represented by chemical formulas.

• Mixtures: Two or more substances physically combined, components retain individual properties, separated by physical means.

  • Homogeneous Mixtures: Uniform properties and composition. A solution is a homogeneous mixture.

  • Heterogeneous Mixtures: Non-uniform composition, component parts distinguishable. Suspensions and colloids are heterogeneous mixtures.

A2.2 Solutions, Suspensions, and Colloids

• Solution: Homogeneous mixture with a solvent (major component) and solute (minor component).

• Suspension: Heterogeneous mixture with visible particles that settle over time and can be separated by filtration.

• Colloid: Heterogeneous mixture with intermediate particle size; particles do not settle and cannot be seen with a microscope.

A2.3 Solubility

• Solubility: Mass of solute that will saturate 100 g of solvent at a given temperature.

• Saturated Solution: Contains as much solute as can be dissolved at a given temperature in the presence of undissolved solute.

• Solubility usually increases with temperature for solid solutes in water.

• Solubility Curve: A graph showing how the solubility of a solute varies with temperature.

A2.4 Separating Mixtures

• Mixtures are separated by physical means; separation method depends on component properties (particle size, boiling point, solubility).

• Filtration: Separates a solid from a liquid using filter paper; solid remains in the filter paper as the residue, the liquid filters through as the filtrate.

• Evaporation: Separates a solid dissolved in a liquid by boiling the solution; the liquid vaporizes as the solute is left behind.

• Crystallization: Separates a solid dissolved in a liquid by allowing the liquid to vaporize slowly at room temperature. The solute remaining has a distinct crystalline structure.

• Simple Distillation: Separates a solution of a solid dissolved in a liquid; the liquid with a lower boiling point vaporizes, cools and condenses back to a liquid at a different location; both the solid and liquid can be collected.

• Fractional Distillation: Separates miscible liquids with close boiling points using a fractionating column.

• Separating Funnel: Separates immiscible liquids with different densities; the denser liquid is drained off.

• Paper Chromatography: Separates dissolved substances based on solubility in a solvent and attraction to the paper.

A2.5 Extraction of Sucrose from Sugar Cane

• Sucrose is extracted from sugar cane using these steps:

  • shred sugarcane;crush and extract juice;

  • clarify;

  • filter;

  • distill;

  • crystallize;

  • centrifuge.

Atomic Structure

A3.1 The Structure of Atoms

• Atom: Smallest component of an element that can exist; the basic building blocks of matter.

• Atoms are made up of protons, neutrons, and electrons (subatomic particles).

• Protons and neutrons are in the nucleus.

  • The number of protons in an atom is known as the atomic number (Z).

• Electrons orbit the nucleus in energy shells.

• Atomic number is unique to each element.

• The number of protons and neutrons in an atom is known as the mass number (A).

• The number of neutrons = A - Z

• Nuclear Notation: ^A_ZX, where X is the atomic symbol.

• Relative charge and mass of subatomic particles.

  • Proton: charge = +1; mass = 1

  • Neutron: charge = 0; mass = 1

  • Electron: charge = -1; mass = \frac{1}{1836}

A3.2 Electronic Configuration of an Atom

• Electrons spin around the nucleus of an atom in a series of levels known as energy shells.

• Each energy shell can hold up to a certain maximum number of electrons:.

  • shell number 1 can hold a maximum of 2 electrons

  • shell number 2 can hold a maximum of 8 electrons

  • shell number 3 may be considered to hold a maximum of 8 electrons.

• The arrangement of electrons in an atom, known as the electronic configuration.

• The electrons in the outermost energy shell are known as valence electrons.

A3.3 Isotopes and Radioactivity

• Isotopes: Different atoms of the same element with the same number of protons and electrons but different numbers of neutrons.

• Isotopes have the same atomic number but different mass numbers.

• Relative Atomic Mass (Ar): Average mass of one atom of an element to one-twelfth the mass of an atom of carbon-12.

• Radioactive isotope: Has an unstable nucleus which decays spontaneously, emitting particles and radiation.

• Radioactive isotopes have many uses, such carbon-14 dating, radiotherapy, radio tracers and energy generation.

The Periodic Table and Periodicity

A4.1 Arrangement of Elements in the Periodic Table

• Periodic table arranges elements in order of increasing atomic number.

• Elements with similar chemical properties are placed in the same group.

• Johann Döbereiner noted triads of elements with similar properties.

• John Newlands proposed the Law of Octaves.

• Dmitri Mendeleev published his 'Periodic Classification of Elements'; arranged based on the element masses and grouped those of similar properties, but it didn't work every time

• Henry Moseley rearranged based on the atomic number instead, which meant other properties aligned along the periodic table as well

• Metals are on the left and Non-metals on the right.

  • Elements on or near this line are metalloids

• The periodic table is divided into 18 groups (vertical columns) and 7 periods (horizontal rows).

A4.2 Trends in Group II of the Periodic Table

• Group II elements are called alkaline earth metals.

• They are metals with similar properties.

• Atoms have 2 valence electrons and form cations with +2 charge.

• Properties change going down the group in increasing reactivity.

A4.3 Trends in Group VII of the Periodic Table

• Group VII elements are called halogens.

• They are non-metals that also have similar properties going down the group.

• Fluorine is the most reactive and astatine is the least reactive.

• Ability to undergo displacement reactions related to oxidising strength. Oxidizing strength decreases going down the group.

A4.4 Trends in Period 3 of the Periodic Table

• Metallic nature decreases and non-metallic nature increases across the period.

• Metallic character decreases from left to the right of the periodic table.

• An electro-negativity increases across the period 3.

Structure and Bonding

A5.1 Principles of Chemical formulas

• Atoms combine with each other to attain the most stable electronic configuration.

• Atoms gain stability if they attain the noble gas configuration of the nearest noble gas to them in the periodic table, i.e. the element that is closest by atomic number.

• Atoms lose, gain or share valence electrons with other atoms.

• A chemical bond involves the loss, gain or sharing the atom's valence electrons.

• There are three main types of chemical bonding: ionic, covalent and metallic.

• There are three main types of chemical formulae: molecular formulae, structural formulae and empirical formulae.

A5.2 Formation Of Ionic Bonds

• Ionic bonding involves the transfer of valence electrons from metal atoms to non-metal atoms

• Ionic compounds result from the strong electrostatic attractive force(s) between both ions

• The ionic compounds have a crystal lattice composed of many of ions held together in a regular, repeating arrangement by ionic bonds

• A cation has positive charge

• An anion has a negative charge

• The sum of the positive charges must equal the sum of the negative charges

A5.3 Writing Chemical Formulae of Ionic Compound

• Metals generally lose electrons forming positive cations, which is named using the element's symbol (e.g., Na^+

• Non-Metals generally gain electrons forming negative anions, named using the element's symbol ended with the suffix "-ide" (e.g., Cl^-

A5.4 Metallic and Covalent Bonds

• Covalent bonding involves the sharing of valence electrons between non-metal atoms.

• Each shared pair of electrons forms a covalent bond

• A molecule is a group of atoms which are bonded together strongly enough to behave as a single units.

• Three of the most common examples of metallic molecules are water, nitrogen, and carbon dioxide.

• Metallic bonding involves the atoms are packed very closely together to form a metal lattice, outer electron shells overlap, and the valence electrons become delocalized