Atoms, Molecules, and Compounds

Atoms, Molecules, and Compounds

• Atoms are the smallest units of matter.
• Atoms consist of subatomic particles.
• The atomic nucleus contains protons and neutrons.
• Protons are positively charged particles; the number of protons defines the atomic number.
• Neutrons are uncharged particles.
• Electrons, which are negatively charged, orbit the nucleus.
• An electron cloud describes the probable locations of electrons.
• A molecule is formed when atoms are physically connected.
• A covalent bond involves the sharing of electrons between atoms.
• A compound occurs when different elements are chemically bonded together.
• An element is a pure substance with atoms that have the same number of protons.
• Chemical elements cannot be broken down into simpler substances by chemical means.
• Calculation of neutrons: number of neutrons = mass number - atomic number.

Elements and Their Properties

• Potassium (K): atomic number 19, 19 protons, 20 neutrons.
• Carbon (C): atomic number 6, 6 protons, 6 neutrons.
• Barium (Ba): atomic number 56, 56 protons, 81 neutrons.
• Chlorine (Cl): atomic number 17, 17 protons, 18 neutrons.
• Argon (Ar): atomic number 18, 18 protons, 21 neutrons.

Common Chemical Compounds

• Examples of chemical compounds used daily are water, sugar, toothpastes, soaps, and baking soda.

The Periodic Table: Groups and Characteristics

• Group 1: Alkali metals are reactive, soft, shiny metals with low melting and boiling points that easily lose one electron to form +1 ions.
• Group 2: Alkaline earth metals are reactive, tend to form +2 ions, and create alkaline hydroxides.
• Group 3: Contains post-transition metals (soft, brittle, weak), metalloids (metallic luster, brittle, semiconductors), and non-metals.
• Group 8: Noble Gases are non-metals with full valence shells, making them unreactive.

Periodic Table Trends

• Atomic number and mass increase from left to right across groups and down through periods.
• Atomic radius increases down groups and decreases from left to right across periods.
• Melting points in periods initially decrease then increase.
• Metal reactivity increases down groups, while non-metal reactivity decreases.
• Metallic character increases down groups and decreases from left to right across periods.

Valence and Electron Configuration

• Valence: The outermost shell of an atom.
• Electron configuration: Summary of electron arrangement around a nucleus.
• Octet Rule: Atoms prefer to have eight electrons in their valence shell (especially elements with atomic numbers up to 20).
• Shell 1 holds up to 2 electrons.
• Atoms in the same group have the same number of valence electrons.
• The number of valence electrons determines an atom's reactivity.
• Shell 2 holds up to 8 electrons.
• Shell 3 holds up to 8 electrons.
• Shell 4 can hold up to 18 electrons.

Covalent vs. Ionic Compounds

• Ionic compounds involve the transfer of electrons and exist as lattices.
• Molecules are usually made of non-metals that share electrons, forming covalent molecular compounds.
• Water, carbon dioxide, oxygen, nitrogen, and carbon tetrachloride are examples of covalent compounds.
• Covalent compounds involve bonds between two non-metals.

Lewis Dot Diagrams

• Lewis Dot diagrams illustrate the arrangement of valence electrons around an atom's symbol.

Formation of Covalent Compounds

• Chlorine (Cl₂): Each chlorine atom shares one electron to achieve a full outer shell.
• Hydrogen Chloride (HCl): Hydrogen and chlorine each share one electron for full outer shells.
• Oxygen (O₂): Each oxygen atom shares two electrons.
• Nitrogen (N₂): Each nitrogen atom shares three electrons.
• Water (H₂O): Each hydrogen needs one electron, and oxygen needs two.
• Carbon Dioxide (CO₂): Each oxygen needs two electrons, while carbon needs four.

Covalent Bonds: Single, Double, and Triple

• Single bond: Involves one pair of electrons (one from each atom).
• Double bond: Involves two pairs of electrons (two from each atom).
• Triple bond: Involves three pairs of electrons (three from each atom).

Chemical Reactions

• Chemical reaction: Substances (reactants) transform through atom rearrangement into new substances (products) with differing properties.
• Signs of a chemical reaction include color change, temperature change, gas formation, precipitate formation, and odor production.
• Endothermic reaction: Absorbs energy (heat) from its surroundings.
• Exothermic reaction: Releases energy (heat or light).

Chemical Reaction Equations

• Ethanol: Ethylene + Water → Ethanol (C₂H₄ + H₂O → CH₃CH₂OH)
• Photosynthesis: Carbon Dioxide + Water → Glucose + Oxygen (6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂)
• Ammonia Boron Trifluoride: Ammonia + Boron Trifluoride → Ammonia Boron Trifluoride (NH₃ + BF₃ → BF₃H₃N)

Chemical Reactions in Non-Living Systems

• Combustion, acid reactions with metals and carbonates, corrosion, precipitation, neutralization, and decomposition.

Reaction Types and Examples

• Combustion: Substance reacts with oxygen, releasing heat and light (e.g., burning wood), is exothermic.
• Acid-Metal Reaction: Acid reacts with a metal to produce hydrogen gas (e.g., acid rain corroding a statue), is exothermic.
• Corrosion: Slow reaction of metals with oxygen and water, forming rust (e.g., rust on iron), is exothermic but slow.
• Precipitation: Mixing two clear solutions forms an insoluble solid (e.g., soap scum in hard water), is neither endothermic nor exothermic.
• Neutralization: Acid reacts with a base to produce water and a salt (e.g., antacid relieves heartburn), is exothermic.
• Decomposition: Single compound breaks down into simpler substances (e.g., hydrogen peroxide breakdown), is endothermic.

Combustion and Energy Production

• Combustion generates electricity in power plants by:
  • Burning fuel to create exhaust gases
  • Exhaust gases spin turbines

Environmental Impacts of Combustion

• Air and noise pollution
• Greenhouse gas emissions contributing to climate change

Alternatives to Combustion-Based Energy

• Renewable energy sources
• Nuclear power
• Hydrogen
• Biomass
• Geothermal energy

Acid-Metal Reactions

• Acids react with metals to produce salt and hydrogen gas.
• Used in metal cleaning and surface treatment.

Corrosion and Infrastructure

• Corrosion damages infrastructure:
  • Weakens materials, leading to potential collapse.
• Prevention:
  • Use non-corrosive materials
  • Keep materials clean and dry

Precipitation and Water Treatment

• Water treatment plants use precipitation to remove contaminants dissolved in a solution.
• Examples: Rain and snow, water filtration systems.
• Importance in chemistry:
  • Filtration process.

Neutralisation and Medicine

• Acidic spills are neutralized with bases; basic spills with acids, creating water and salt.
• pH balance is essential for bodily functions like digestion and oxygen delivery.

Decomposition and Waste Management

• Decomposition releases nitrogen into the atmosphere.
• In composting and landfill management, it transforms organic matter into compost and provides soil nutrients.

Problem-Solving Scenarios

• Neutralizing Acid Spill: Use basic substances like sodium carbonate to neutralize, and the reaction forms water and salt, protecting the ecosystem.
• Firefighting: Combustion requires fuel, oxygen, and heat, remove these to control a fire.
  • CO₂ extinguishers replace oxygen reducing heat.
• Water is not effective on flammable liquids (Class B), electrical (Class C), or metal (Class D) fires.

Ionic Compounds

• Ions are formed when atoms lose or gain electrons.
• Ionic compounds are formed when electrons are transferred.
• Dependent on number of electrons in outer shall
• Noble gases are non-reactive due to already having a full shell.
• Chemical reactions are caused by atoms wanting a full outer shell.
• If loses 3 electrons - write +3 at top - Na^+3
• If gains 2 electrons - write - 2 at the top
• Non-metal ions are always negative
• Non-metal ions end in -ide
• Cations: positively charged ions + attracted to electrons
• Anions: negatively charged ions
• Ionic compounds found in lattices and dont exist as molecules.
• In a lattice - all ions are in an orderly arrangement.

Questions Section

• A negative charge (Br-) gains 1.
• Ionic compound example is Magnesium Sulfide.

Formation Example

• Magnesium flouride.
• Magnesium - Mg2+ - is a cation as it loses 2 electrons.
• Flouride - F- - is an anion as it gains one electron.

Transfer/Eagles

• Silicon (Group 4): 4 valence electrons, may gain or lose 4 electrons (Si+4, Si-4).
• Fluorine (Group 7): 7 valence electrons, gains 1 electron (F-1).
• Magnesium (Group 2): 2 valence electrons, loses 2 electrons (Mg+2).

Compound Creation

• Magnesium Fluoride (MgF₂): Magnesium loses 2 electrons to 2 Fluorine atoms.
• Magnesium Silicide (Mg₂Si): Two Magnesium atoms lose 2 electrons each to Silicon.

Neutralisation Experiment

• Aim: investigate between sodium carbonate (Na₂CO₃) and hydrochloric acid (HCl).
• Chemical Equation: Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂
• Word Equation: Sodium carbonate + Hydrochloric acid → Sodium chloride + Water + Carbon dioxide
• Colour from universal indicator changed from red to green to purple which meant neutralisation had occurred
• A candle went out when lit near test tube, proof of gas evolution
  • Hydrochloric indicator - red/orange
  • Neutralisation starts - yellow/light green
  • Complete neutralisation - green
  • Excess sodium carbonate - blue/purple

Conclusion

• Salt, water and carbon dioxide are formed when acid is fully reacted.
• Double diplacemnt reaction

Different Types of Reactions

• Combustion Reactions general equation:
  • Fuel + Oxygen → Carbon dioxide + Water + Energy
  • CH₄ + 2O₂ → CO₂ + 2H₂O + Energy
• Extinguisher removed on of the elements.
• Remove oxygen ; fire blanket or foam extinguisher
• Remove heat ; water or fire extinguisher
• Fuel ; shut off gas or fuel source and remove any flammables

Precipitation Reactions

• Two reactions to form a sediment
• AgNO₃ + NaCl → AgCl↓ + NaNO₃
• solid product

Corrosion Reactions

• Gradual reaction with there environment.
• Rust and frosty later observed.
• Iron Oxide
• Iron + Oxygen + Water → Hydrated Iron(III) Oxide (Rust)
• Galvinization and coat.

Decomposition Reaction Practical

• Neutralise: to stop from having an effect.
• Acid and base counteract.
• When reacted salt and water are produced.
• Hydrochloric acid and sodium hydroxide:
  • NaOH (aq) + HCl (aq) → H₂O(I) + NaCl(aq)

Neutralisation Examples

• Ingredients of antacid tablets neutralize the acid without producing carbon dioxide gas.
• AgNO₃(aq) + NaCl(aq) → AgCl (s) + NaNO₃(aq)
• Metal reacts with oxygen, water, acids or other chemicals
• prevent rust by not having the material come into contact with corrosive substances.
• decomposition of zinc carbonate represented by:
  • ZnCO3(s) → ZnO (s) + CO2 (g)
• Burning methane:
  • CH4(g) +202(g) → CO2 (g) + 2H2O (g)
  • methane + oxygen → carbon dioxide + water