Chemistry Notes PDF

1. Metallic Bonding Structure of Metals: Metals are made up of closely packed positive ions arranged in a giant lattice structure. The outer electrons of each metal atom are delocalized, forming a "sea" of free-moving electrons. The electrostatic attraction between these delocalized electro...

1. Metallic Bonding Structure of Metals: Metals are made up of closely packed positive ions arranged in a giant lattice structure. The outer electrons of each metal atom are delocalized, forming a "sea" of free-moving electrons. The electrostatic attraction between these delocalized electrons and the positive metal ions is the metallic bond. Detailed Explanation of Properties: 1. High Melting and Boiling Points: The strong electrostatic forces between positive ions and delocalized electrons require a significant amount of energy to overcome. Metals like tungsten have extremely high melting points due to strong metallic bonds. 2. Electrical Conductivity: Delocalized electrons move freely, carrying an electric current. Example: Copper is used in electrical wiring due to its excellent conductivity. 3. Thermal Conductivity: Heat energy is transferred through the movement of delocalized electrons and vibrations of ions. 4. Malleability and Ductility: When metals are bent or hammered, layers of ions slide over one another while the delocalized electrons adjust to maintain bonding. 2. Basics of Atoms and Elements The Atom: Composed of three main subatomic particles: Protons: Positive charge, found in the nucleus. Neutrons: No charge, found in the nucleus. Electrons: Negative charge, orbiting the nucleus in shells. Key Atomic Concepts: Atomic Number: Number of protons in the nucleus. Mass Number: Total number of protons and neutrons. Isotopes: Atoms with the same number of protons but different numbers of neutrons. Example: Carbon Carbon-12: 6 protons, 6 neutrons, 6 electrons. Carbon-14: 6 protons, 8 neutrons, 6 electrons. 3. States of Matter Particle Arrangement and Properties: State Particle Arrangement Movement Energy Properties Solid Tightly packed in fixed positions Vibrate in place Low Fixed shape and volume Liquid Close but not fixed Slide past each other Moderate Fixed volume, takes shape of container Gas Far apart and random Move rapidly and freely High No fixed shape or volume Changes of State: Melting: Solid to liquid; particles gain energy and break free from fixed positions. Boiling: Liquid to gas; particles gain enough energy to overcome attractions. Freezing: Liquid to solid; particles lose energy and settle into fixed positions. Condensation: Gas to liquid; particles lose energy and form closer bonds. Sublimation: Solid to gas directly (e.g., dry ice). 4. The Periodic Table (Continued) 3. Group 0 (Noble Gases): These elements (e.g., Helium, Neon, Argon) have a full outer electron shell and are thus chemically inert(unreactive). Properties: Low boiling points that increase down the group (due to stronger intermolecular forces). Non-flammable and used in light bulbs (argon) or balloons (helium). 5. Types of Bonding Ionic Bonding (not fully covered yet but here’s a preview): Occurs between metals and non-metals. Metal atoms lose electrons, becoming positively charged cations. Non-metal atoms gain electrons, becoming negatively charged anions. The oppositely charged ions are held together by strong electrostatic forces in a lattice structure. Example: Sodium chloride (NaCl). Covalent Bonding: Occurs between non-metals. Atoms share pairs of electrons to achieve a full outer shell. Example: Water (H₂O) – Oxygen shares electrons with two hydrogen atoms. 6. Chemical Reactions Signs of a Chemical Reaction: Color change. Temperature change (heat given off or absorbed). Formation of gas (bubbles). Formation of a precipitate (solid). Conservation of Mass: In a chemical reaction, the total mass of reactants equals the total mass of products (mass is conserved). Example:
2
H
2

+
O
2

→
2
H
2

O


2H2 +O2 →2H2 O
Mass of hydrogen and oxygen reactants equals the mass of water produced. 7. Elements, Compounds, and Mixtures Element: A pure substance made of only one type of atom. Example: Oxygen (O₂), Gold (Au). Compound: A substance made of two or more elements chemically combined in fixed proportions. Example: Carbon dioxide (CO₂). Mixture: A combination of two or more substances not chemically bonded and can be separated by physical methods. Example: Air (a mixture of gases). 8. Separation Techniques 1. Filtration: Used to separate an insoluble solid from a liquid. Example: Sand and water. 2. Distillation: Separates liquids based on differences in boiling points. Example: Separating water from saltwater. 3. Chromatography: Separates substances based on their solubility and movement through a medium (e.g., paper). Example: Separating pigments in ink. 4. Evaporation and Crystallization: Evaporation removes the liquid, leaving dissolved substances to form crystals. Example: Obtaining salt from seawater. 9. Acids, Bases, and pH (Preview) Acids: Substances that release hydrogen ions (H⁺) in water. Example: Hydrochloric acid (HCl). Bases: Substances that release hydroxide ions (OH⁻) in water or neutralize acids. Example: Sodium hydroxide (NaOH). pH Scale: Measures the acidity or alkalinity of a solution (0-14). Acids: pH < 7, Bases: pH > 7, Neutral: pH = 7 (e.g., water). 10. Chemical Equations and Balancing 1. What is a Chemical Equation? A chemical equation represents a chemical reaction using symbols and formulas. It shows the reactants (substances you start with) and products (substances formed). Example:
Reactants
→
Products


Reactants→Products
2
H
2

+
O
2

→
2
H
2

O


2H2 +O2 →2H2 O 2. Steps to Write a Balanced Chemical Equation Balancing ensures the Law of Conservation of Mass is followed (mass of reactants = mass of products). Steps to Balance: 1. Write the unbalanced equation:
Example:
H
2

+
O
2

→
H
2

O


H2 +O2 →H2 O 2. Count the atoms of each element on both sides of the equation: Left side: 2 H, 2 O Right side: 2 H, 1 O 3. Adjust coefficients to balance atoms: Add a coefficient of 2 before H₂O:
2
H
2

+
O
2

→
2
H
2

O


2H2 +O2 →2H2 O 4. Double-check to ensure all elements are balanced: Left side: 4 H, 2 O Right side: 4 H, 2 O 3. Examples of Balancing 1. Magnesium and Oxygen Reaction:
Unbalanced:

M
g
+
O
2

→
M
g
O


Unbalanced:Mg+O2 →MgO
Balanced:

2
M
g
+
O
2

→
2
M
g
O


Balanced:2Mg+O2 →2MgO 2. Combustion of Methane:
Unbalanced:

C
H
4

+
O
2

→
C
O
2

+
H
2

O


Unbalanced:CH4 +O2 →CO2 +H2 O
Balanced:

C
H
4

+
2
O
2

→
C
O
2

+
2
H
2

O


Balanced:CH4 +2O2 →CO2 +2H2 O 11. Ionic vs. Covalent Compounds 1. Ionic Compounds Formed between metals and non-metals. Metals lose electrons to form positive ions (cations). Non-metals gain electrons to form negative ions (anions). Oppositely charged ions are held together by strong electrostatic forces (ionic bonds). Properties of Ionic Compounds: 1. High melting and boiling points: Strong ionic bonds require lots of energy to break. 2. Conduct electricity in molten or dissolved states: Ions are free to move. 3. Soluble in water: Many ionic compounds dissolve easily. Examples: Sodium chloride (NaCl). Magnesium oxide (MgO). 2. Covalent Compounds Formed between non-metals. Atoms share pairs of electrons to achieve a full outer shell. Molecules are held together by covalent bonds. Properties of Covalent Compounds: 1. Low melting and boiling points: Weak intermolecular forces (between molecules). 2. Do not conduct electricity: No free-moving charges. 3. Often insoluble in water: Exceptions include polar covalent compounds like sugar. Examples: Water (H₂O). Carbon dioxide (CO₂). Methane (CH₄). Comparison of Ionic and Covalent Compounds: Property Ionic Compounds Covalent Compounds Bond Type Transfer of electrons (electrostatic forces) Sharing of electrons (covalent bonds) Melting/Boiling Points High Low Electrical Conductivity Conduct in molten/dissolved state Do not conduct Solubility in Water Generally soluble Generally insoluble (except polar)