Chemical Bonding Ms. Manning

Back to Compounds….. 2 Types: Covalent Compounds Formed when non-metals bond with other non- metals Ionic Compounds Formed when metals bond with non-metals

Conductivity in Liquid Classification Compound of: Bonding Structure Tm and Tb Conductivity in Solid Conductivity in Liquid Metals Metallic Lattice High Metals and Non Metals Ionic Low Non Metals Covalent Molecular

Properties of Metallic compounds Relatively dense solids (exception Hg) Good conductors of heat and electricity Lustrous when clean/ freshly cut Strong, malleable (can be shaped) and ductile (can be drawn into a wire) Sonorous: Ringing sound when hit Relatively high melting and boiling points Usually form positive ions

Properties of Non-Metals Non-lustrous Can exist in any state - generally gases at room temperature Brittle, non-ductile Poor conductors of heat and electricity Usually exist as molecules in their elemental form Low densities, melting and boiling points. Combine with other nonmetals to form covalent bonds Generally form negative ions, e.g.  Cl-, SO42-, and N3-

Properties of Metalloids Generally look metallic but are brittle (not malleable or ductile) Neither good conductors or insulators; instead they are semiconductors.

Chemical Bonding Chemical Bond = the force of attraction holding atoms or ions together This is how compounds are made!

Classifying Compounds Ionic Compound = a pure substance formed from a metal and a nonmetal NaCl CaO Molecular Compound = a pure substance formed from two or more different nonmetals SO2 CO2

Ionic versus Molecular Compound Electrical Conductivity = the ability of a material to allow electricity to flow through it Ionic Compounds  conduct electricity Molecular Compounds  DO NOT

Electrolyte Electrolyte = a substance that forms a solution that conducts electricity Ionic compounds  form electrolytic solutions Molecular compounds  form non- electrolytic solutions

Ionic Bonding Ions = atoms that have gained or lost electrons Ionic Bond = the electrostatic attraction between positive and negative ions in a compound Metals lose electrons Non-metals gain electrons Both form octets = MORE STABLE

Ionic Bonding – Bohr Diagrams

Lewis Dot Diagrams – Ionic Bonding KBr MgCl2

Naming Ionic Bonding Ionic Compounds Metal + Non-metal Metal name  same as on the atom name Non-Metal  suffix “-ide” Example: NaCl = Sodium Chloride LiF = Lithium Flouride MgO = Magnesium Oxide

Non-metal Suffixes Nitrogen = Nitride Oxygen = Oxide Fluorine = Fluoride Phosphorus = Phosphide Sulfur = Sulfide Chlorine = Chloride Selenium = Selenide Bromine = Bromide Iodine = Iodide Text: page 73 Q. 9 -13

How Many Atoms in a Molecule? Diatomic Molecules = a molecule consisting of two atoms of the same or different elements CO Polyatomic Molecules = a molecule consisting of more than two atoms of the same or different elements NH3

Covalent Bonding Covalent Bond = the attractive forces between two atoms that results when electrons are shared by the atoms A simultaneous attraction of two nuclei for a shared pair of electrons In Lewis Diagrams – the shared pairs of electrons are shown as lines and the lone pairs as dots

Octet Rule Still Applies! The shared pair of electrons is considered to be a pair of electrons that make both atoms have an octet 8 8

Lewis Dot Diagrams – Covalent Bonds

The Lone Pair Lone Pair = a pair of valence electrons not involved in bonding

Bonding Capacity Bonding Capacity = the number of electrons lost, gained or shared by an atom when it bonds chemically Allows us to predict how many bonds an atom can form

Bonding Capacity Carbon 4 Nitrogen 5 3 Oxygen 6 2 Halogens 7 1 Atom # of Valence Electrons Number of Bonding Electrons Bonding Capacity Carbon 4 Nitrogen 5 3 Oxygen 6 2 Halogens 7 1 Hydrogen

Choosing the Central Atom for Polyatomic Molecules The central position… Is usually occupied by the element with the highest bonding capacity C and N are often in the central position The least electronegative atom is usually the central atom Hydrogen is NEVER the central atom Oxygen and Halogens are usually not the central atom Page 79 gives step by step instructions

Covalent Bonds = Strong A large amount of energy is needed to separate the atoms that make up molecules The stronger the bond the greater the amount of energy needed to break the bond Single bond = strong Double bond = stronger Triple bond = strongest

Single, Double & Triple Bonds

Polar Covalent Bonds Polar Covalent Bonds = a covalent bond formed between atoms with significantly different electronegativities; a bond with some ionic characteristics When electrons are shared between two atoms = covalent bond In a bond between identical atoms the electrons are shared equally In a bond between two different atoms the sharing is unequal

Non-Polar versus Polar Covalent

Comparison…

Difference in Electronegativity… If the electronegativity difference is: less than 0.2 = bond is pure covalent is between 0.2 and 1.6 = bond is polar covalent is greater than 1.7 = bond is ionic

Polar Molecules Polar Molecules = a molecule that is slightly positively charges at one end and slightly negatively charged at the other because of electronegativity differences

Types of Forces Intramolecular Force = the attractive forces between atoms and ions within a compound Ionic Polar Covalent Non-polar Covalent Intermolecular Force = the attractive force between molecules

IntRA versus IntER-molecular Forces

Some Intermolecular Forces 3 major types of Intermolecular Forces: Dipole-dipole forces London dispersion forces Hydrogen bonding The first two are known as van der Waals Forces London dispersion forces and dipole-dipole forces

van der Waals Force Dipole-dipole force = an attractive force acting between polar molecules Attraction between oppositely charged ends of polar molecules

van der Waals Force London Dispersion force = an attractive force acting between all molecules including nonpolar molecules A result of temporary displacements of the electron “cloud” around the atoms in a molecule resulting in a extremely short- lived dipole

London Dispersion Force

Hydrogen Bonding Hydrogen Bonding = a relatively strong dipole-dipole force between a positive hydrogen atom of one molecule and a highly elecgtronegative atom (F, O or N) in another molecule

Hydrogen Bonding

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