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Chemistry A: Elements, Compounds, and Mixtures

1. ElementsDefinition:Classification of Elements:Periodic Table Trends:2. CompoundsDefinition:Types of Compounds:Chemical Bonding:3. MixturesDefinition:Types of Mixtures:Separation Techniques:4. Differences Between Elements, Compounds, and Mixtures5. States of Matter and ChangesStates of Matter:Changes in State:6. Purity and FormulationsMeasuring Purity:7. Empirical and Molecular FormulaEmpirical Formula:Molecular Formula:Calculation of Empirical Formula:8. Transition Metals and Coordination ComplexesTransition Metals:Coordination Complexes:

1. Elements

Definition:

  • A pure substance consisting of only one type of atom.

  • Cannot be broken down into simpler substances by chemical means.

Classification of Elements:

  • Metals: Good conductors of heat & electricity, malleable, ductile, high melting/boiling points (e.g., Fe, Cu, Al).

  • Non-metals: Poor conductors, brittle, lower melting/boiling points (e.g., O, N, S).

  • Metalloids: Properties intermediate between metals and nonmetals (e.g., Si, B).

  • Noble Gases: Inert, unreactive due to a full valence electron shell (e.g., He, Ne, Ar).

  • Halogens: Highly reactive nonmetals, form salts with metals (e.g., Cl, Br, I).

Periodic Table Trends:

  • Atomic Radius: Decreases across a period, increases down a group.

  • Ionization Energy: Increases across a period, decreases down a group.

  • Electronegativity: Increases across a period, decreases down a group.

2. Compounds

Definition:

  • A substance formed when two or more elements are chemically bonded together in fixed proportions.

Types of Compounds:

  • Ionic Compounds: Metal + Non-metal (e.g., NaCl, MgO)

    • High melting/boiling points, conduct electricity in molten or aqueous state.

    • Form crystal lattice structures.

  • Covalent Compounds: Non-metal + Non-metal (e.g., CO₂, H₂O)

    • Low melting/boiling points, do not conduct electricity.

    • Exist as molecules.

  • Metallic Compounds: Metal + Metal (e.g., Alloys like brass, steel)

    • Malleable, ductile, conducts electricity and heat.

Chemical Bonding:

  • Ionic Bonding: Transfer of electrons from metals to non-metals.

  • Covalent Bonding: Sharing of electron pairs between atoms.

  • Metallic Bonding: Sea of delocalized electrons around metal cations.

3. Mixtures

Definition:

  • A combination of two or more substances physically combined, not chemically bonded.

  • Components retain their individual properties.

Types of Mixtures:

  • Homogeneous Mixture: Uniform composition (e.g., saltwater, air).

  • Heterogeneous Mixture: Non-uniform composition (e.g., sand and water, salad dressing).

Separation Techniques:

  • Filtration: Separates solids from liquids (e.g., sand from water).

  • Distillation: Separates liquids based on boiling points (e.g., ethanol from water).

  • Chromatography: Separates based on solubility and adsorption (e.g., ink components).

  • Evaporation: Removes a liquid to leave a solid residue (e.g., salt from seawater).

  • Magnetism: Separates magnetic materials from non-magnetic ones (e.g., iron filings from sulfur powder).

4. Differences Between Elements, Compounds, and Mixtures

Property

Element

Compound

Mixture

Composition

Single type of atom

Two or more atoms chemically bonded

Two or more substances physically combined

Separation

Cannot be broken down

Chemical methods required

Physical methods

Properties

Unique to the element

Different from individual elements

Retain properties of individual components

Example

Oxygen (O₂)

Water (H₂O)

Air (N₂, O₂, CO₂)

5. States of Matter and Changes

States of Matter:

  • Solid: Definite shape and volume, strong intermolecular forces.

  • Liquid: Indefinite shape, definite volume, moderate intermolecular forces.

  • Gas: Indefinite shape and volume, weak intermolecular forces.

Changes in State:

Process

Description

Example

Melting

Solid → Liquid

Ice to Water

Freezing

Liquid → Solid

Water to Ice

Evaporation

Liquid → Gas

Water to Steam

Condensation

Gas → Liquid

Steam to Water

Sublimation

Solid → Gas

Dry Ice (CO₂)

Deposition

Gas → Solid

Frost formation

6. Purity and Formulations

  • Pure Substance: Contains only one type of element or compound.

  • Impure Substance: Contains multiple elements or compounds.

  • Formulation: A mixture designed for a specific purpose (e.g., medicine, alloys, fertilizers).

Measuring Purity:

  • Boiling/Melting Point Analysis: Pure substances have fixed boiling/melting points, while impure substances show variations.

  • Chromatography: Used to identify components in a mixture.

7. Empirical and Molecular Formula

Empirical Formula:

  • Simplest whole-number ratio of atoms in a compound.

  • Example: CH₂O (for glucose C₆H₁₂O₆).

Molecular Formula:

  • Actual number of atoms of each element in a molecule.

  • Example: C₆H₁₂O₆ (glucose).

Calculation of Empirical Formula:

  1. Find mass (or %) of each element.

  2. Convert mass to moles.

  3. Divide by the smallest number of moles.

  4. Write an empirical formula.

8. Transition Metals and Coordination Complexes

Transition Metals:

  • Found in the d-block of the periodic table.

  • Exhibit variable oxidation states.

  • Form colored compounds due to d-orbital electron transitions.

  • Act as catalysts in many chemical reactions.

  • Examples: Fe (iron) in hemoglobin, Cu (copper) in electrical wiring, Cr (chromium) in stainless steel.

Coordination Complexes:

  • Consist of a central metal ion surrounded by ligands.

  • Ligands donate electron pairs to the metal (e.g., H₂O, NH₃, Cl⁻).

  • Coordination number: Number of ligand attachment points.

  • Common complex ions:

    • [Cu(NH₃)₄]²⁺ (tetraamminecopper(II)) – Deep blue color.

    • [Fe(CN)₆]³⁻ (hexacyanoferrate(III)) – Yellow color.

    • [Ag(NH₃)₂]⁺ (diammine silver(I)) – Used in Tollens' reagent.

  • Applications: Catalysis, biological processes (e.g., hemoglobin with Fe²⁺), industrial dyes, and metal plating.


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Chemistry A: Elements, Compounds, and Mixtures

1. Elements

Definition:

  • A pure substance consisting of only one type of atom.

  • Cannot be broken down into simpler substances by chemical means.

Classification of Elements:

  • Metals: Good conductors of heat & electricity, malleable, ductile, high melting/boiling points (e.g., Fe, Cu, Al).

  • Non-metals: Poor conductors, brittle, lower melting/boiling points (e.g., O, N, S).

  • Metalloids: Properties intermediate between metals and nonmetals (e.g., Si, B).

  • Noble Gases: Inert, unreactive due to a full valence electron shell (e.g., He, Ne, Ar).

  • Halogens: Highly reactive nonmetals, form salts with metals (e.g., Cl, Br, I).

Periodic Table Trends:

  • Atomic Radius: Decreases across a period, increases down a group.

  • Ionization Energy: Increases across a period, decreases down a group.

  • Electronegativity: Increases across a period, decreases down a group.

2. Compounds

Definition:

  • A substance formed when two or more elements are chemically bonded together in fixed proportions.

Types of Compounds:

  • Ionic Compounds: Metal + Non-metal (e.g., NaCl, MgO)

    • High melting/boiling points, conduct electricity in molten or aqueous state.

    • Form crystal lattice structures.

  • Covalent Compounds: Non-metal + Non-metal (e.g., CO₂, H₂O)

    • Low melting/boiling points, do not conduct electricity.

    • Exist as molecules.

  • Metallic Compounds: Metal + Metal (e.g., Alloys like brass, steel)

    • Malleable, ductile, conducts electricity and heat.

Chemical Bonding:

  • Ionic Bonding: Transfer of electrons from metals to non-metals.

  • Covalent Bonding: Sharing of electron pairs between atoms.

  • Metallic Bonding: Sea of delocalized electrons around metal cations.

3. Mixtures

Definition:

  • A combination of two or more substances physically combined, not chemically bonded.

  • Components retain their individual properties.

Types of Mixtures:

  • Homogeneous Mixture: Uniform composition (e.g., saltwater, air).

  • Heterogeneous Mixture: Non-uniform composition (e.g., sand and water, salad dressing).

Separation Techniques:

  • Filtration: Separates solids from liquids (e.g., sand from water).

  • Distillation: Separates liquids based on boiling points (e.g., ethanol from water).

  • Chromatography: Separates based on solubility and adsorption (e.g., ink components).

  • Evaporation: Removes a liquid to leave a solid residue (e.g., salt from seawater).

  • Magnetism: Separates magnetic materials from non-magnetic ones (e.g., iron filings from sulfur powder).

4. Differences Between Elements, Compounds, and Mixtures

Property

Element

Compound

Mixture

Composition

Single type of atom

Two or more atoms chemically bonded

Two or more substances physically combined

Separation

Cannot be broken down

Chemical methods required

Physical methods

Properties

Unique to the element

Different from individual elements

Retain properties of individual components

Example

Oxygen (O₂)

Water (H₂O)

Air (N₂, O₂, CO₂)

5. States of Matter and Changes

States of Matter:

  • Solid: Definite shape and volume, strong intermolecular forces.

  • Liquid: Indefinite shape, definite volume, moderate intermolecular forces.

  • Gas: Indefinite shape and volume, weak intermolecular forces.

Changes in State:

Process

Description

Example

Melting

Solid → Liquid

Ice to Water

Freezing

Liquid → Solid

Water to Ice

Evaporation

Liquid → Gas

Water to Steam

Condensation

Gas → Liquid

Steam to Water

Sublimation

Solid → Gas

Dry Ice (CO₂)

Deposition

Gas → Solid

Frost formation

6. Purity and Formulations

  • Pure Substance: Contains only one type of element or compound.

  • Impure Substance: Contains multiple elements or compounds.

  • Formulation: A mixture designed for a specific purpose (e.g., medicine, alloys, fertilizers).

Measuring Purity:

  • Boiling/Melting Point Analysis: Pure substances have fixed boiling/melting points, while impure substances show variations.

  • Chromatography: Used to identify components in a mixture.

7. Empirical and Molecular Formula

Empirical Formula:

  • Simplest whole-number ratio of atoms in a compound.

  • Example: CH₂O (for glucose C₆H₁₂O₆).

Molecular Formula:

  • Actual number of atoms of each element in a molecule.

  • Example: C₆H₁₂O₆ (glucose).

Calculation of Empirical Formula:

  1. Find mass (or %) of each element.

  2. Convert mass to moles.

  3. Divide by the smallest number of moles.

  4. Write an empirical formula.

8. Transition Metals and Coordination Complexes

Transition Metals:

  • Found in the d-block of the periodic table.

  • Exhibit variable oxidation states.

  • Form colored compounds due to d-orbital electron transitions.

  • Act as catalysts in many chemical reactions.

  • Examples: Fe (iron) in hemoglobin, Cu (copper) in electrical wiring, Cr (chromium) in stainless steel.

Coordination Complexes:

  • Consist of a central metal ion surrounded by ligands.

  • Ligands donate electron pairs to the metal (e.g., H₂O, NH₃, Cl⁻).

  • Coordination number: Number of ligand attachment points.

  • Common complex ions:

    • [Cu(NH₃)₄]²⁺ (tetraamminecopper(II)) – Deep blue color.

    • [Fe(CN)₆]³⁻ (hexacyanoferrate(III)) – Yellow color.

    • [Ag(NH₃)₂]⁺ (diammine silver(I)) – Used in Tollens' reagent.

  • Applications: Catalysis, biological processes (e.g., hemoglobin with Fe²⁺), industrial dyes, and metal plating.