Chemical Compounds & Bonds

• A compound is a pure substance
• It is made by a fixed combination of elements
• Elements are held together by chemical bonds

Chemical Bonding

• Deals with the Octet Rule and the behavior of electrons
• There are three types of elements

Types of elements

• Metals tend to lose electrons in chemical reactions
• Non-metals tend to gain or share electrons in chemical reactions
• Semi-metals (metalloids) tend to share electrons in chemical reactions

Bond Theory

• An atom can fill its outer (valence) shell in multiple ways
• Electrons can be gained, lost, or shared
• Bonds are classified based on electron activity
• Valence refers to the bonding capacity of an atom or ion

Metallic bonds

• Form between metal atoms
• Characterized by “pooled” electrons

Ionic Bonds

• Typically form between metals and non-metals
• They involve "transferred" electrons, creating cations and anions

Covalent Bonds

• Form between non-metals (and semi-metals)
• Characterized by "shared" electrons

Ionic Bonding

• Sodium has 1 valence shell electron
• Sodium loses one electron to form a monovalent cation
• Losing 1 valence shell electron creates a positive charge
• Chlorine has 7 valence shell electrons
• Chlorine gains one electron to form a monovalent anion
• Gaining 1 electron forms a negative charge

Covalent Bonding

• A covalent bond involves atoms sharing electrons to fill their outer shell
• Covalent bonding occurs when atoms are close enough for valence electrons to be attracted to both nuclei

Bonds by number of electrons shared

• A bond with two shared electrons is a single covalent bond
• A single covalent bond has two total electrons shared
• A bond with 4 shared electrons is a double covalent bond, with four total electrons shared
• A bond with 6 shared electrons is a triple covalent bond, with six total electrons shared
• Atoms may use more than one bond to fulfill their octet, hydrogen needs only a duet

Types of Compounds

• Metallic compounds contain metal atoms held together by metallic bonds
• Ionic compounds contain at least one ionic bond in the formula unit
• Ionic bonds typically occur between metals and non-metals, such as sodium chloride (NaCl)
• Molecular compounds contain only covalent bonds, made of non-metals, like water (H2O)

Metallic Compounds

• Composed of all metal atoms combined by an "electron sea" of delocalized valence shell electrons
• Metals tend to lose electrons, forming cations
• Lost electrons act as a "glue," holding the metal cations together

Ionic Compounds

• Composed of cations and anions bound together by at least one ionic bond
• Cations are usually a metal and anions are usually one or metals
• The basic unit of an ionic compound is the formula unit, the smallest, electrically neutral collection of ions

Example Ionic Compound

• Table salt (NaCl) is composed of Na+ and Cl- ions in a one-to-one ratio

Molecular Compounds

• Contain only covalent bonds and only non-metals (or semi-metals)
• Water is composed of H2O molecules
• Propane is composed of C3H8 molecules

Molecular Formula

• Molecular Formula expresses the composition of a chemical compound
• Indicates the actual number of each type of atom in a molecule
• Butane (C4H10) has 4 carbon atoms and 10 hydrogen atoms
• Aluminum chloride (AlCl3) has 1 aluminum and 3 chlorine atoms

Formulas

• C4H10 (the actual # of atoms of each element) = molecular formula
• CH3‒CH2‒CH2‒CH3 = condensed formula

Formula Mass

• The mass of an individual molecule or formula unit (reported in amu)
• The sum of the masses of the atoms in a single molecule or formula unit
• C4H10 has a formula mass of (4 C * 12.01) + (10 H * 1.01) = 58.14 amu
• AlCl3 has a formula mass of (1 Al * 26.98) + (3 Cl * 35.45) = 133.33 amu

Atomic Mass Unit (amu)

• The mass of a single proton is ~1.66 x 10-24 g
• The atomic number (Z) is the number of protons
• The mass number (A) is the number of protons and neutrons
• The atomic mass is an average mass of all natural isotopes
• A single atom of carbon-12 is 1.99 x 10-23 g, or (12 x 1.66 x 10-24 g) or 12 AMU
• A single atom of carbon weighs 12.01 amu on average

Coulomb’s Law

• Describes the magnitude of the electrostatic force of attraction or repulsion between two charged particles
• The force is directly proportional to the product of the magnitudes of charges
• The force is inversely proportional to the square of the distance between the charges

Periodic Trend: Atomic Radius

• Atomic radius is the distance from the nucleus to the valence shell electrons
• Atomic radius increases moving down a column because electrons are organized in shells
• The valence shell is farther from the nucleus as you move down a column
• Atomic radius decreases moving from left to right across a period
• The nuclear charge increases moving left to right

Ionic Bond Length/Strength

• Two properties determine the strength of an ionic bond
• Larger charges result in stronger and shorter ionic bonds
• The distance between the two charged particles
• Smaller radii result in stronger and shorter ionic bonds

Covalent Bond Length/Strength

• The distance between the nuclei of covalently bonded atoms
• Smaller atoms have shorter bond lengths
• Shorter bonds are stronger and longer covalent bonds are weaker
• Single bonds tend to be longer and weaker than double bonds
• Triple bonds are shorter and stronger than double bonds

Electronegativity

• Is the attraction for shared electrons
• Electronegativity increases across a period (left to right) and decreases down a group (top to bottom)
• Noble gases are not assigned Electronegativity values, because they do not share electrons
• Fluorine is the most electronegative element

Ionic Compounds

• Composed of one or more cations and one or more anions that are bound together by at least one ionic bond
• The cation (metal) is always listed first when writing or naming ionic compounds
• An ionic bond is an electrostatic (charge-based) interaction

Properties of Ionic Compounds

• Often contain one or more metal cations and one or more non-metal anions
• Metals tend to lose electrons to form positively charged cations
• Non-metals tend to gain electrons to form negatively charged anions
• Bound together by at least one ionic bond
• There are some ions (called polyatomic ions) that also contain covalent bonds
• Are solids at room temperature because ionic bonds are very strong
• Contain charged particles (ions), so they conduct electricity when dissolved in water

Electrolytes

• Is a substance which conducts an electric current when dissolved in water
• A soluble ionic compound is an electrolyte
• A strong electrolyte is a completely soluble ionic compound
• A weak electrolyte is partially soluble ionic compound
• A non-electrolyte is a molecular compound or an insoluble ionic compound
• An aqueous solution is when a solution is dissolved in water

Molecular Compounds

• They contain only covalent bonds and only non-metals (or semi-metals)
• Generally list carbon first, then hydrogen, then any other atom in alphabetical order when writing or naming
• A directional bond is a covalent bond with electrons in place between the two atoms
• Do not contain charged particles, so they do not conduct electricity
• Molecular compounds are non-electrolytes

Acids

• Are an exception to molecular compounds being non-electrolytes
• Acids are a molecular compounds (all covalent bonds) that ionize when dissolved in water
• Acids are written with hydrogen listed first

Molecular Mass

• Physical state of molecular compounds at room temperature is partly related to their molecular mass

Molecular Mass & Physical State

• Gases: methane (CH4 = 16.05 amu), carbon dioxide (CO2 = 44.01 amu)
• Liquids: acetic acid (C2H4O2 = 60.06 amu), isopropyl alcohol (C3H8O = 60.11 amu)
• Solids: glucose (C6H12O6 = 180.18 amu)
• Notable exceptions: water (H2O, 18.02 amu) is a liquid at room temperature

Ionic vs. Molecular Compounds

• There is a continuum between ionic and covalent bonds

Polar Covalent Bond

• Has unequally shared electrons

Bond Polarity

• It is caused by differences in electronegativity
• A polar covalent bond still involves shared electrons, but they are not shared evenly
• The electron spends more time nearer to the more electronegative atom
• Partial positive and negative charges are a result
• Partial negative charge on the more electronegative atom
• Partial positive charge on the less electronegative atom

Polar Covalent Bonds as Vectors

• Magnitude of the vector is the difference in electronegativity
• The vector points toward the more electronegative atom
• The bond between hydrogen and fluorine would be a polar covalent bond
• Hydrogen's electronegativity is 2.1, and fluorine's electronegativity is 4.0

Determining Ionic vs. Covalent Bond

• If it is a metal and a non-metal, it is an ionic bond
• If both atoms bonded are non-metals (or semi-metals) it is a covalent bond
• Determine the difference in electronegativity (ΔΕneg) of the bonded atoms

Electronegativity Difference

• If ΔΕneg > 0.4, it is a polar covalent bond
• If ΔΕneg = 0 to 0.4, it is a non-polar covalent bond

Covalent Bond Vectors

• In covalent bonds, to determine if a covalent bond is polar, assess the difference in electronegativity (ΔΕneg)
• Covalent bonds occur with (non-metals & semimetals).
• If the difference is > 0.4, there is a polar bond.
• If ΔΕneg is 0 to 0.4, it is non-polar.
• To determine if a covalent bond is polar, determine the difference in electronegativity (ΔΕneg) of the bonded atoms
• If the difference is > 0.4, it is a polar bond
• If ΔΕneg is 0 to 0.4, it is non-polar

Polarity

• For a molecule to be polar, there are two requirements
• There must be at least 1 polar bond, and, the molecule must have a non-symmetrical shape
• A bond is polar if there are differences in electronegativity

Molecular Shape

• Molecules have a 3-dimensional shape, and covalent bonds have a “bond angle.”
• There are 3 basic shapes for all molecular compounds, and they depend on the number of electron domains present on an atom.
• The atom in the middle is called the central atom
• An electron domain is an atom bonded to the central atom denoted with an “X” or a non- bonding electron pair

Symmetry

• To a chemist, symmetry means that all of the electron domains are identical

Non-Polar Molecules

• Carbon Dioxide is non-polar

Polar Molecules

• Carbon oxysulfide is polar

The formula for Boron Trichloride

• BCl3, is trigonal planar with a bond angles of 120 degrees

Ammonia

• Has 4 electron domains, has a trigonal pyramidal shape- non-symmetrical and is polar

Water

• Has 4 electron domains and a bent shape- non-symmetrical and is polar

Symmetrical Molecules

• CH4 and CCl4 are symmetrical; the other 3 molecules do not have identical electron domains so they are not symmetrical

Water (H2O) Properties

• Water, made of only 3 atoms, contains 10 protons, 10 electrons (and ~8 neutrons)
• Each water molecule weighs only 18.02 amu
• At normal atmospheric pressure, water is a liquid at room temperatures or ~25 ºC
• compared to other molecules
  • nitrogen (N2) weighing 28.02 amu at -196 ºC
  • oxygen (O2) weighing 32.00 amu at -183 ºC
  • carbon dioxide (C02) is 44.01 amu at -78 ºC
• Water is polar, a good solvent, and biologically important

Hydrogen Bonding

• Water properties results from hydrogen bonding from polar molecules in hydrogen which are covalently bonded to N, O, or F
• This forms strong interactions called hydrogen bonds (H-bonds)
• A hydrogen covalently bonded to N, O, or F forms an H- bond with an N, O, or F on an adjacent molecule
• A hydrogen bond is a physical, not a chemical, change

Hydrogen Bonding & Electronegativity

• It arises in part from the high electronegativity of nitrogen, oxygen, and fluorine
• When hydrogen is bonded to very electronegative elements, hydrogen’s nucleus is exposed
• Water loving compounds contain functional groups with –OH and/or ‒NH2
• These can form hydrogen bonds
• Hydrogen bonding also lowers boiling points

Electrolytes

• An electrolyte is a substance which, when dissolved in water, conducts an electric current
• An aqueous solution contains water
• Organic chemistry studies carbon-based compounds called hydrocarbons
• Valence refers to the bonding capacity of an atom or ion, and the electrons in the outer-most shell are called valence shell electrons

Four Valence Shell Electrons

• Carbon has four valence shell electrons
• the Octet Rule says that atoms like to have 8 valence Electrons.
• Carbon will form 4 bonds to get an octet

Hydrocarbons

• Are non-polar compounds, as they only have C-C and C-H bonds, which are non-polar
• As nonpolar, they cannot form hydrogen bonds
• Tend to have lower boiling points that similarly sized compounds that can form H-bonds
• Also have lower heat capacities (they can absorb less heat) than polar compounds
• Exhibit lower values for other physical properties

Fuel Types

• Alkane contains only single bonds between the carbons
• Alkene contains 1 or more C=C double bonds
• Alkyne contains 1 or more C≡C triple bond

Energy Types

• During combustion the process of combining a fuel with molecular oxygen
• Potential energy is stored energy, or the energy of position, or of chemical bonds
• Kinetic energy is the energy of motion, or ofreaction

Fossil Fuels

• Are non-renewable combustible organic substances
• they will be used up one day
• Their waste products have adverse effects on the environment

Combustion

• Bonds store the energy to hold them together and it releases them when they are consumed
• CCoulomb’s Law predicts that shorter bonds tend to be stronger bonds
• Combustion reactions produce energy typically in the form of heat and light flame
• An example reaction is that 1 CH4 (g) + 2 O2 (g) → 1 CO2 (g) + 2 H2O (g)