Unit 7: Bonding

Why Bond? Elements bond in order to get a stable valence e- configuration (a full valence shell) Noble gases have stable valence e- configurations and tend not to bond Bonding is all about becoming more stable (at the lowest energy possible)

Breaking/Forming of Bonds When a bond is broken, energy is ABSORBED (required) Ex) F2 (g) + ENERGY  F (g) + F (g) When a bond is formed, energy is RELEASED (given off) Ex) F(g) + F(g)  F2 (g) + ENERGY

Types of Chemical Bonds Chemical bonds are formed when valence electrons are: Transferred from one atom to another (ionic) Shared between atoms (covalent) Mobile within a metal (metallic)

Don’t forget about electronegativity! The ability of an atom to attract electrons when in a compound F- 4.0 (highest EN) and Fr-0.7 (lowest EN) Noble gases- no assigned value EN values can be used to determine the type of bond that forms between two atoms Bigger difference in EN ( difference >1.7 )  more ionic character

Ionic- Metal to Non-Metal An ionic bond is the force of attraction that holds ions of opposite charge together, forming an ionic compound ***Electrons are transferred from M NM ***Metal loses electrons (becomes positive) and the non-metal gains electrons (becomes negative) Large difference in electronegativity between the metal and non-metal Overall charge of compound is neutral

Writing Lewis Dot Structures for Ionic Compounds sodium and chlorine barium and oxygen calcium and fluorine potassium and sulfur

Covalent Bonds A bond between two nonmetals that involves a sharing of electrons (“tug of war”) Can have EQUAL or UNEQUAL sharing Small or no difference in electronegativity Overall charge is neutral Covalent compounds are called molecules

Types of Covalent Bonds Non-Polar Covalent Equal sharing of electrons Polar covalent Unequal sharing of electrons Coordinate Covalent One nonmetal provides BOTH e- to be shared

Non-Polar Covalent Bond Equal sharing of e- Look for two of the same nonmetals All diatomics have non-polar bonds! Remember: H2 O2 F2 Br2 I2 N2 Cl2 Ex) F2

Polar Covalent Bond Unequal sharing of electrons Different NMs with different electronegativity values Ex) HF The more electronegative atom has a slight (-) charge The less electronegative atom has a slight (+) charge

Bond Polarity Bond Polarity: refers to a separation of charge in a bond Ex) HCl : partial (+) charge on H and partial (-) charge on Cl This separation of charge is called a dipole ***The greater the difference in electronegativity, the greater the polarity. Ex) HF is more polar than HBr

Formation of Covalent Bonds

Formation of Covalent Bonds Bond length: the average distance between two bonded atoms Bond energy: the energy required to break a chemical bond and form neutral atoms (kJ/mol) In general, bond energies become larger as bond lengths become shorter

Drawing Lewis Structures for Covalent Compounds After drawing your diagram, all atoms MUST have 8 valence e- (except for hydrogen, which should have 2 valence e-) When elements from Group 14 are involved in a covalent bond, they spread their e- out Carbon tends to form 4 covalent bonds Atoms of carbon can bond to each other to form chains, rings and networks

Lewis Structures may contain… Single covalent bond: one pair of e- is shared between two atoms (2 total e-) Double covalent bond: 2 pairs of e- are shared between atoms (4 total e-) Triple covalent bond: 3 pairs of e- are shared between atoms (6 total e-)

Coordinate Covalent Bond The third type of covalent bond… One non-metal donates BOTH electrons to be shared Examples: NH3 and H2O

PROPERTIES OF IONIC AND COVALENT COMPOUNDS

Ionic Compounds An ionic compound exists as a collection of positively and negatively charged ions arranged in repeating patterns A formula unit is the lowest whole-number ratio of ions in an ionic compound NaCl (1:1 ratio) , MgCl2 (1:2 ratio)

Properties of Ionic Compounds (salts) Ionic bonds are very strong bonds, so all are solids (hard with a crystalline structure) High melting points and boiling points Soluble (can dissolve) in water Cannot conduct electricity as a solid (ions are not free to move) Can conduct electricity as a liquid or in an aqueous solution (ions are free to move) Electrolytes- substances that conduct electricity when dissolved in water

Properties of Covalent Compounds (Molecules) Many are gases or liquids at room temperature Molecular solids tend to be soft Low melting points and boiling points Cannot conduct electricity in any phase Generally insoluble in water Except sugars! (C12H22O11)

Network Covalent Solids A special case of covalent bonding All atoms are held together in a very strong covalent network Ex. Carbon (Diamond) and SiO2 Properties: Very hard High m.p. and b.p. Poor conductors

Metallic Bonding: The “sea of mobile electrons”

Metallic Bonding Holds metal atoms together Electrons are mobile and can move from one atom to another, creating (+) charged metal ions Charged metal ions are immersed in a “sea of mobile electrons”

IMFs vs. Bonds

LD Forces

Dipole-Dipole Attractions

Hydrogen bonding

London Dispersion Forces London Dispersion forces: weakest of all molecular interactions; caused by temporary shifts in charge Between nonpolar molecules http://antoine.frostburg.edu/chem/senese/101/liquids/faq/h-bonding-vs- london-forces.shtml The bigger the atom or molecule  the greater the strength of dispersion forces  the higher the BP

Dipole-Dipole Interactions Dipole interactions: attraction between polar molecules The positive and negative charges of different molecules attract each other Ex. HCl http://chemmovies.unl.edu/ChemAnime/DIPOLED/DIPOLED.html

Hydrogen bonds: A special case of dipole-dipole interactions Hydrogen bonds: intermolecular force between the H of one molecule and a highly electronegative atom of another molecule (must be N, O, or F) Ex. H2O, NH3 ***The high b.p. of water is due to hydrogen bonding