IB Chem: Bonding & Structure

Ionic Model, Covalent Model, Metallic Model & From Models to Materials

Lattice Dissociation Enthalpy

• the standard enthalpy change that occurs on the formation of 1 mole of gaseous ions form the solid lattice

• endothermic process

Cations

positively charged ions

Ions

• An electrically charged atom formed by the loss or gain of electrons

Anion

• negatively charged ion

Cation

• positively charged ion

Ionic Bond

• strong electrostatic force of attraction between oppositely charged ions

Giant Ionic

• high melting and boiling points

• not volatile

  • Volatility: the vaporization of a chemical

• Soluble — can form ion-dipole bonds

• only conduct electricity when molten or in solution

• hard, brittle

• electrostatic attraction between ions

• at room temp - solid

  • e.g. NaCl

Giant Metallic

• moderately high melting/boiling point

• electrical conductivity only when solid or liquid

• insoluble

• hard, malleable

• solid — at room temp

• delocalised electrons attracting positive ions

  • e.g. copper

Simple Covalent

• low melting/boiling point

• does not conduct electricity

• usually insoluble unless polar

• soft

• at room temp - solid, liquid or gas

• weak, intermolecular forces and covalent bonds within a molecule

  • e.g. Br2

Giant Covalent

• very high melting/boiling point

• does not conduct electricity

  • (except graphite)

• insoluble

• very hard(diamond and silicone) or soft(graphite)

• electrons in covalent bonds between atoms

  • e.g. graphite, silicone(IV), oxide

Metallic Bonding

• the structure of metallic bonding has positive metal ions suspended in a “sea” of delocalised electrons

  • very strong electrostatic forces between the positive metal centres

Delocalised Electrons

free moving electrons not bound to their atom

Malleability

• to be malleable —> when a force is applied, the metal layers can slide

  • e.g. metals can be hammered into sheets

• attractive forces between the metal ions and electrons act in all directions

Strength

metallic compounds are strong and hard due to the strong attractive forces between the metal ions and delocalised electrons

Electrical Conductivity

• can conduct electricity in the solid/liquid state

  • due to mobile electrons that freely move around and conduct electricity

• outer electrons increase across a period, # of delocalised charges increases

  • Na = 1 outer electron

  • Mg = 2 outer electrons

  • Al = 3 outer electrons

    • ability to conduct electricity increases across a period

Ductile

• the ability for a metal to be stretched or drawn into a thin wire without breaking

Thermal Conductivity

• metals are good thermal conductors due to the behaviour of their cations and delocalised electrons

  • cations in metal lattic vibrate vigorously as their thermal energy increases

  • delocalised electrons carry increased kinetic energy and transfer it rapidly throughout the metal —contributes to high thermal conductivity

Melting & Boiling Point

• metals have high melting/boiling points

• due to strong electrostatic forces of attraction between the cations and delocalised electrons in the metallic lattice

Alloy

a mixture of two or more metals or nonmetals — to produce a substance with metallic properties

• mixed physically but not chemically combined

• ions of the different metals are spread throughout the lattice and are bound together by the delocalized electrons

Uses of metals

• Aluminum is used in food cans because it is non-toxic and resistant to corrosion and acidic food stuffs

• copper is used in electrical wiring b/c it is a good electrical conductor and malleable/ductile

• stainless steel is used for cutlery as it is strong and resistant to corrosion

Melting point trends

• melting point of metal increases moving across a period, from left to right

  • greater chard difference leads to a stronger electrostatic attraction, and therefore a stronger metallic bond

Ionization Energy

The amount of energy required to remove an electron from an isolated atom/molecule

Electron Affinity

how readily a neutral atom(in gaseous state) will attract and hold onto an extra electron, forming a negative ion

• the measure of an atom’s “attractiveness” for an electron

Electronegativity

the ability of an atom to draw an electron pair towards itself in a covalent bond

Trends in Melting Points of Metals

strength of electrostatic attraction can be increased by:

• increasing the number of electrons

• increasing the number of positive charges

• decreasing the size of the cations(ionic radius)

Brass

• substitutional alloy

• copper & zinc

• strong & resistant to corrosion

  • e.g door handles, hinges

Steel

• interstitial alloy

• iron, carbon and other elements

• very strong

  • e.g. construction, bridges, cars

Stainless steel

• iron, chromium, nickel & carbon

• corrosion resistant

  • cutlery, surgical instruments, cookware

Solder

• lead and tin

• low melting point

  • e.g., joining metals in electrical circuits metals and jewelry

Bronze

• copper tin

• hard and strong resistance to corrosion

• medals & sculptures

Linear

• e.g. BeCl2, CO2 & ethyne

• angle = 180 degrees

• two electron domains

• 0 lone pairs

Triangular Planar/Trigonal Planar

• 3 electron domains

• 0 lone pairs

• 120 degrees

• e.g. BF3, CH2O

  • boron trifluoride, ethene and methanal

Molecular Geometry: Bent Linear

• Domain geometry: Trigonal Planar

• e.g. SO2(sulfure dioxide)

• 118 degrees

• 1 lone pair

• ‘expands the octet’

Tetrahedral

• 109.5 degrees

• 0 lone pairs

  • e.g. methane & ammonium(CH4 & NH4+)

Molecular Geometry: Trigonal Pyramidal

• Domain Geometry: Tetrahedral

• 107 degrees

• 1 lone pair

  • e.g. ammonia(NH3)

Molecular Geometry: Bent/Angular Linear

• Domain Geometry: Tetrahedral

• 104.5 degrees

• 2 lone pairs

  • e.g. water

Van der Waals’ forces

• a term used to include:

  • london dispersion forces

  • dipole- induced dipole attractions

  • dipole-dipole attractions

• forces occur between molecular(intermolecularly) as well as within moelcule(intramolecularly)

London dispersion forces

• instantaneous induced dipole - induced dipole forces that exist between all atoms and molecules

• caused by temporary dipoles

  • constantly appearing and disappearing due to constant motion of electrons

• weakest bond

Strength depends:

• the number of electrons in the molecule

• surface area of the molecules

Hydrogen Bonding

• strongest type of intermolecular force

  • special type of permanent dipole - permanent dipole bonding

Trigonal Bypyramidal

• composed of a central atom and five surrounding atomms

• Bond angles: 90, 120, 180

  • e.g. PCl5

Seesaw

• 3-D

• central atom has one lone pair & 4 bonding pairs

• Bond angles < 180

  • 90 - 120

  • e.g. SF4

VSPER Theory

electron pairs(both bonding and non-bonding) repel each other and arrange themselves to minimize repulsion