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7. Bonding
Ionic Bonding
Definition: A type of chemical bonding characterized by the transfer of electrons from one atom to another.
Components:
Metal: Typically loses electrons and becomes positively charged ions (cations).
Non-Metal: Typically gains electrons and becomes negatively charged ions (anions).
Key Processes:
Electrons are transferred: The metal gives up one or more electrons to the non-metal.
Formation of Ions: After electron transfer, the atoms form ions with opposite charges that attract each other, resulting in an ionic bond.
Covalent Bonding
Definition: A type of chemical bonding characterized by the sharing of electrons between atoms.
Components:
Non-Metal + Non-Metal: Covalent bonds typically occur between non-metals.
Key Processes:
Electrons are shared: Each atom contributes one or more electrons to be shared in order to achieve a full outer shell.
Skills Required
Draw Lewis Diagrams
Ability to visually represent the bonding in compounds through Lewis dot structures.
Identify Type of Bonding
Determine whether a bond is ionic, covalent, or another type based on the elements involved.
Explain How Bonding Works
Articulate the principles and processes behind ionic and covalent bonding.
8. Periodic Trends & Reasoning
Reactivity
Focus on: Alkali metals and halogens.
Alkali metals (Group 1) are highly reactive, especially with water.
Halogens (Group 17) are reactive non-metals that readily form compounds with alkali metals.
Conceptual Understanding:
Explain the rationale behind why elements gain or lose electrons based on their atomic structure.
Connection to Valence Electrons: The number of valence electrons determines an element's reactivity.
Skills Required
Explain Why: Must be able to give detailed reasons for the observed reactivity trends and not just identify them.
9. Naming Ionic and Covalent Compounds
Naming Ionic Compounds
Rule: Cations are named first, followed by anions.
Example: Sodium chloride from sodium (Na+) and chloride (Cl-).
Naming Covalent Compounds
Method: Use prefixes to indicate the number of atoms.
Prefixes: Mono-, di-, tri-, tetra-, etc.
Writing Formulas
Ability to write chemical formulas for both ionic and covalent compounds based on names.
Example: H2O for water (two hydrogen atoms and one oxygen atom).
4. The Periodic Table
Groups vs Periods
Groups: Vertical columns that share similar chemical properties.
Periods: Horizontal rows that represent energy levels.
Classification of Elements
Metals: Generally conductors of electricity and heat, malleable, and ductile.
Non-Metals: Typically poor conductors and brittle.
Metalloids: Exhibit properties of both metals and non-metals.
Families of the Periodic Table
Alkali Metals (Group 1)
Highly reactive, especially with water.
Alkaline Earth Metals (Group 2)
Reactive but less than alkali metals.
Halogens (Group 17)
Very reactive non-metals.
Noble Gases (Group 18)
Generally inert due to full valence shells.
Concepts to Understand
Reactivity Trends: Understand how reactivity changes across periods and down groups.
Valence Electrons: The outermost electrons that dictate bonding behavior.
Why Noble Gases are Stable: Complete valence shell results in lack of reactivity.
5. Bohr-Rutherford & Lewis Diagrams
Bohr Diagrams
Task: Draw Bohr diagrams that accurately represent the arrangement of electrons in shells.
Lewis Dot Diagrams
Task: Draw Lewis dot diagrams to denote valence electrons.
Identify Valence Electrons
The focus is on the outermost electrons relevant for bonding.
6. Ions
Definition of Ions
Ion: An atom or molecule that has lost or gained one or more electrons, resulting in a net charge.
Types of Ions
Cations (+): Positively charged ions formed when an atom loses electrons.
Anions (-): Negatively charged ions formed when an atom gains electrons.
Process of Gaining/Losing Electrons
Atoms gain or lose electrons to achieve a stable electron configuration, often resembling the nearest noble gas configuration.
Charges of Common Groups
Group 1 (Alkali Metals): Often form +1 cations.
Group 2 (Alkaline Earth Metals): Generally form +2 cations.
Group 17 (Halogens): Typically form -1 anions.
Group 16: Often form -2 anions.
Skills Required
Determine Ion Charge: Calculate the charge of ions based on the number of electrons gained or lost.
Draw Ions: Accurate drawings of ions can be represented using both Bohr and Lewis diagrams.
Matter & Changes
Definition of Matter
Matter: Anything that has mass and occupies space.
States of Matter:
Solid: Definite shape and volume.
Liquid: Define volume but takes the shape of its container.
Gas: Neither definite shape nor volume.
Particle Theory
Key Ideas:
All matter is made up of particles.
Particles are in co nstant motion and interact with one another.
Temperature affects the energy and motion of particles.
Changes in Matter
Physical Changes: Changes that do not alter the chemical composition of a substance.
Example: Melting ice into water.
Chemical Changes: Changes that result in the formation of new substances.
Evidence of Chemical Change (5 Signs):
Change in color.
Production of gas.
Formation of precipitation.
Change in temperature.
Change in properties.
Physical vs Chemical Properties
Physical Properties: Characteristics that can be observed without changing the substance's composition (e.g., color, melting point).
Chemical Properties: Characteristics that determine how a substance reacts with other substances (e.g., flammability).
Classification of Matter
Element: A pure substance that cannot be broken down into simpler substances.
Compound: A substance made of two or more elements chemically combined.
Mixture: A combination of two or more substances that retain their individual properties.
Homogeneous: Uniform composition (e.g., saltwater).
Heterogeneous: Non-uniform composition (e.g., salad).
Skills Required
Identifying examples of each category: element, compound, and mixture.
Atomic Structure
Fundamental Particles
Protons: Positively charged particles located in the nucleus.
Neutrons: Neutral particles also located in the nucleus.
Electrons: Negatively charged particles that orbit the nucleus in electron clouds.
Atomic Number vs Mass Number
Atomic Number: The number of protons in the nucleus, which determines the element.
Mass Number: The sum of protons and neutrons in the nucleus.
Calculation of Subatomic Particles
Calculating Protons: Equal to the atomic number.
Calculating Neutrons: .
Calculating Electrons: Equal to the number of protons in a neutral atom.
Isotopes
Definition: Variants of the same element that have the same number of protons but differ in the number of neutrons.
Example: Carbon-12 and Carbon-14 are isotopes of carbon, differing in their neutrons.