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Formula Unit
The simplest ratio of cations (positive ions) to anions (negative ions) in an ionic compound.
Monatomic Ion
An ion consisting of only one atom (e.g., Na⁺, Cl⁻, O²⁻).
Polyatomic Ion
An ion consisting of multiple atoms bonded together with an overall charge (e.g., NH₄⁺, SO₄²⁻).
Oxyanion
A polyatomic ion containing oxygen (e.g., NO₃⁻, SO₃²⁻).
Transition Metals
Can have multiple oxidation states (e.g., Fe²⁺ and Fe³⁺).
Binary Ionic Compounds
Compounds formed from a metal and a nonmetal.
Naming Binary Ionic Compounds
Name the metal first, then the nonmetal with '-ide' ending.
Naming Transition Metals
Use a Roman numeral in parentheses to indicate charge.
Naming Polyatomic Ionic Compounds
Name the cation first, then the polyatomic ion, without changing the ending.
Crystalline Structures
Ionic compounds form these structures, with the formula unit representing the entire neutral crystal.
Balancing Charges
Ensure the total charge in the formula equals zero.
Polyatomic Ions in Formulas
Function as a single unit; use parentheses with a subscript if multiple copies are needed.
Polyatomic ions
More than one atom acting as a single unit.
Naming rules
Cation first, anion second; '-ide' for monatomic anions; use Roman numerals for transition metals; polyatomic ions keep their name.
Writing formulas
Balance charges to get a neutral compound; use parentheses for multiple polyatomic ions.
Metallic bond
The electrostatic attraction between positively charged metal cations and a 'sea' of delocalized valence electrons.
Characteristics of a metallic bond
Allows metals to have unique properties such as conductivity, malleability, and ductility.
Electron sea model
A model explaining metallic bonding where valence electrons move freely among metal cations.
Delocalized electrons
Electrons that are not bound to any specific atom and can move freely within a metallic solid.
Alloy
A mixture of elements that retains metallic properties.
Malleability
Metals can be hammered into thin sheets without breaking.
Electric Potential
When an electric potential (voltage) is applied, electrons flow freely, allowing metals to conduct electricity.
Luster
The interaction of delocalized electrons with light causes metals to be shiny (lustrous) as they absorb and re-emit photons.
Hardness and Strength
Metals with more delocalized electrons are harder and stronger.
Properties of Alloys
Alloys have enhanced properties compared to pure metals, such as greater strength, resistance to corrosion, and improved hardness.
Types of Alloys
Two types of alloys: Substitutional alloy → Atoms of similar size replace each other; Interstitial alloy → Smaller atoms fill gaps between larger atoms.
Ionic Compounds vs Metals
Ionic compounds: rigid lattice, brittle, electrons fixed in place; Metals: electron sea model, malleable, delocalized electrons.
Metallic Bonding and Conductivity
Electrical conductivity: Delocalized electrons move freely; High boiling point: Strong attraction between metal cations and the electron sea requires a lot of energy to break.
Ionic vs Metallic Bonding
Ionic bonding: Electrostatic attraction between oppositely charged ions; Metallic bonding: Attraction between metal cations and a sea of delocalized electrons.
Ionic Bond
The electrostatic force holding oppositely charged ions together.
Ionic Compound
A chemical compound formed from ionic bonds.
Electrolyte
A substance whose aqueous solution conducts electricity.
Crystal Lattice
A 3D geometric arrangement of ions in an ionic compound.
Lattice Energy
Energy needed to separate 1 mole of ions in a crystal lattice.
Formation of Ionic Bonds
Ions form bonds by transferring electrons from metal to nonmetal, creating oppositely charged ions that attract each other.
Exothermic Reactionterm-17
Ionic bond formation is exothermic because energy is released when the oppositely charged ions attract and form a stable compound.
Exothermic Reaction
Ionic bond formation releases energy, making the compound more stable.
Electron Transfer Process Example 2
Calcium loses 2 electrons to become Ca²⁺, and each fluorine gains 1 electron to become F⁻, resulting in CaF₂.
Binary Ionic Compounds
Contain only two different elements: a metallic cation and a nonmetallic anion.
Sodium Nitride
Na₃N, formed when sodium (Na) loses 1 electron and nitrogen (N) gains 3 electrons.
Strength of Ionic Bonds
Ionic bonds are very strong, leading to high melting and boiling points, hardness, brittleness, and the ability to conduct electricity when dissolved in water or melted.
Example of Crystal Lattice
In NaCl, each Na⁺ ion is surrounded by 6 Cl⁻ ions, forming a cubic crystal.
Electrical Conductivity in Solid State
No conductivity (ions are fixed in place).
Electrical Conductivity in Molten or Dissolved State
Ions are free to move, making the compound an electrolyte.
Factors Affecting Lattice Energy
Ion Size: Smaller ions have higher lattice energy because the charges are closer together.
Ion Charge and Lattice Energy
Higher charges increase lattice energy.
Example of Ion Charge
MgO (Mg²⁺ and O²⁻) has higher lattice energy than NaF (Na⁺ and F⁻).
Ionic Compound Neutrality
The total positive charge of the cations equals the total negative charge of the anions, balancing the compound.
Energy During Ionic Bond Formation
Energy is released (exothermic), leading to a lower energy, more stable structure.
Conduct Electricity in Solution
Free ions can move and carry current when dissolved.
Lattice Energy and Ionic Bond Strength
Higher lattice energy means stronger ionic bonds due to either smaller ions or higher ion charges.
Chemical Bond
The force that holds two atoms together.
Ionic Bonds
Formed by the attraction between positive ions (cations) and negative ions (anions).
Covalent Bonds
Formed by the sharing of electrons between atoms.
Cation
A positively charged ion formed by losing electrons.
Anion
A negatively charged ion formed by gaining electrons.
Ionization Energy
The energy needed to remove an electron from an atom.
High Ionization Energy
Harder to remove an electron (e.g., noble gases).
Low Ionization Energy
Easier to remove an electron (e.g., alkali metals).
Electron Affinity
The energy change when an atom gains an electron.
High Electron Affinity
Atom readily gains electrons (e.g., halogens).
Low Electron Affinity
Atom doesn't easily gain electrons (e.g., noble gases).
Naming Anions
Add "-ide" to the element's name (e.g., Chlorine → Chloride (Cl⁻)).
Cation Formation
Requires input of energy (ionization energy).
Anion Formation
Releases energy (electron affinity).
Attraction in Chemical Bonds
Attraction between positive nucleus of one atom and negative electrons of another; attraction between positive ions (cations) and negative ions (anions).
