Chapter 8
5
Penta
any acid that contains hydrogen and 1 other nonmetal
Binary acid
Composed only of two nonmetals atoms; covalent bond
Binary molecular compounds
1. Occurs due to random movement of electrons 2. If the electrons are more on one side of an atom than another it produces an instantaneous dipole (temporary) 3. This will then induce a dipole in the atom next to it 4. Weak forces because they don't last long (weakest of the Intermolecular Forces) 5. Occurs between nonpolar substances and noble gases
(London) Dispersion Forces (a Van der Waals force)
Ate
-ic acid
Ite
-ous acid
Hydrogen atom consists of _ proton and _ electron. H+ is just 1 proton.
1
P4, S8
2 common molecules
N2, O2, F2, Cl2, Br2, I2, H2
7 diatomic molecules
- A = central atom - X = ligand (an atom or group of atoms) - E = lone pair on the central atom * helps identify electron geometry
AXE Notation
Substances that ionize (break up into + or - ions) in aqueous (aq; dissolved in water) solutions to form hydrogen ions (H^1+)
Acids
1. Depends on how strongly each of the bonded atoms attracts electrons 2. Steps: - determine the electronegativity of the individual atoms bonded - compared the difference to the scale - electronegativity difference shows a character of relationship; it is often not clearly ionic or covalent
Bond Character
1. Symmetric molecules are usually "nonpolar" molecules 2. Asymmetric molecules are always "polar" as long as the bong type is "polar"
Bond Polarity and Molecular Shape
- generally as bond length increases, bond-dissociation energy decreases - as the number of bonds increases, bond-dissociation energy increases
Bond energy
- The distance between the nuclei of two bonded at the positive of maximum attraction - the length vibrates (oscillates) between the two atoms 1. If the atoms pull apart: the (+) and (-) will attract and bring them in 2. if the atoms get too close: the (+) and (+) will repel and spread them out 3. When the net (overall) attraction is greater than the net repulsion, the atoms bond covalently and a molecule is formed - At an ideal length, the molecule will have a low potential energy, therefore a high stability
Bond length
- the greater the # of shared pairs, the shorter the bond length - the fewer the # of shared pairs, the longer the bond length
Bond length determined by the number of electron pairs the atoms share
- the larger the atoms, the longer the bond length - the smaller the atoms, the shorter the bond length
Bond length determined by the size of the two bonding atoms
Amount of energy required to break a specific covalent bond (or form)
Bond-dissociation energy
Localized region where bonding electrons will most likely be found
Bonding orbital
- combinations of non-metals - contains covalent bonds (share electrons)
Covalent compounds
10
Deca
2
Di
1. Polar molecule have ends that are oppositely charged 2. In a collection of polar molecules, the dipoles tend to orient themselves so that the (+) dipole on one molecule is near the (-) dipole on another molecule 3. The more polar the molecule, the stronger the dipole force 4. The farther apart, the weaker the dipole froce
Dipole-dipole forces (a Van der Waal force)
- shares 2 pairs of electrons - shorter bond length than single bonds - higher bond energy than single bonds - mostly occurs with the elements C, O, N, S
Double covalent bond
1. Determine the type and # of atoms in the molecule 2. Determine the total number of valence electrons in the atoms to be combined. ADD electrons if the charge of the polyatomic ion is negative. SUBTRACT electrons if the charge is positive 3. Arrange the atoms to form a skeleton structure 4. Add unshared pairs of electrons so that each atom has 8, except H (only needs 2) 5. Count the electrons in the structure to be sure that the number of valence electrons used equals the number available 6. "Place the whole structure in [ ] with a superscript equal to the charge outside of the square brackets
Drawing Lewis Structures of Polyatomic Ions
1. Is a measure of the tendency of an atom to accept an electron 2. Increases across the period, decreases down the group 3. Values are measured experimentally
Electron Affinity
1. Indicates the relative ability of an atom to attract electrons in a chemical bond 2. Increases across the period, decreases down the group 3. Values were assigned by Linus Pauling - F given the value of 3.98 or 4
Electronegativity
1. Angle = 180˚ 2. 2 terminal atoms
Linear
1. Enthalpy is the heat of content of a system at constant pressure (how much energy in a substance) - (change of H) is + value for breaking bonds - (change of H) is - value for forming bonds 2. Enthalpy of reaction (change of Hrxn) is the change of enthalpy for a reaction - enthalpy of reaction = sigma (sum) (enthalpy of bonds broken) + (enthalpy of bonds formed) - endothermic reaction is + - exothermic reaction is -
Enthalpies of Reaction
Hydrocyanic acid - HCN(aq)
Exception to naming acids
1. Incomplete Octet: if the molecule has an odd number of valence electrons - lose 1 electron from the central atom (weaker atom) 2. Suboctet: if the central atoms contains fewer than eight valence electrons - since the suboctet is unstable and the atom doesn't have an electron to share, a coordinate covalent bond will form - a coordinate covalent bond is the bond formed when 1 nonmetal atom donates both of its electrons to be shared with an atom or ion that needs two electrons to from a stable electron configuration with a lower potential energy 3. Expanded OctetL if the central atom contains more than eight valence electrons - can occur when the central atom is located on the 3, 4, 5, 6, or 7 energy level
Exceptions to the Octet Rule
- Occurs when the particles within one atom attract and repel particles within the other atom. When these forces balance, a bond forms
Formation of a covalent bond
- group 17 forms 1 covalent bond with an atom of a nonmetal - group 16 forms 2 covalent bonds with atoms of nonmetals - group 15 forms 3 covalent bonds with atoms of nonmetals - group 14 forms 4 covalent bonds with atoms of nonmetals
Groups and single bonds
7
Hepta
6
Hexa
1. H is always on the outside 2. Halogens (7 valence electrons) are usually on the outside or end of the molecule. Those on the end are called terminal atoms. 3. When a molecule contains more atoms of one kind than the other, the kind that there is the most of tends to surround the kind that there is the least of 4. The atom with the smallest electronegativity is often the central atom
Hints to drawing Lewis Structures
1. Process in which atomic orbitals mix and from new, identical hybrid orbitals 2. Example with Carbon: - Carbon has only two half-filled orbitals and should form only 2 bonds with 2 hydrogen atoms. However, the stable compound formed from carbon and hydrogen is CH4 (methane) which has 4 bonds And 4 hydrogen atoms - The explanation is the "s and p" orbitals combine to form a new hybrid orbital called "sp^3" because it formed from 1s and 3p orbitals 3. The number of hybrid orbitals formed equals the number of atomic orbitals mixed
Hybridization Theory/Model
1. The attraction occurring when a H atom bonded to a strongly electronegative atom such as F, O, or N, is also attracted to another strongly electronegative atom of a different molecule 2. H is a strong (+) dipole. It will attract to the (-) dipole of another molecule
Hydrogen Bonding
1. Attractions BETWEEN the molecules 2. Determine the physical properties 3. Three types - hydrogen bonding - dipole-dipole forces (a Van der Waal force) - (London) Dispersion Forces (a Van der Waals force)
Intermolecular Forces
1. Attractions WITHIN the molecule 2. Determine chemical properties (what does it react with) 3. Two types: - ionic bonds from the transfer of electron(s) - covalent bonds from the sharing of electron(s)
Intramolecular Forces
1.7 - 4.0
Ionic bond
1. a diagram showing the arrangement of valence electrons among the atoms 2. Valence electrons (electrons present in the outermost energy level of the atom 3. Are used to represent covalent bonds 4. Each atom has a stable octet (8) of electrons 5. Unshared pairs (lone pairs) is a pair of electrons that is not involved in covalent bonding. It belongs to only one atom 6. A shared pair can be represented with a dash (2 electrons) 7. Represents chemical formula - symbols represent nucleus and core electrons - the dot pairs or dashes between the atomic symbols represent electrons pairs in covalent bonds - a pair of dots adjacent to only 1 atomic symbol represent a lone pair
Lewis Structure
EG: Linear
MG: Linear
EG: Octahedral
MG: Octahedral, Square Pyramidal, Square Planar
EG: Tetrahedral
MG: Tetrahedral, trigonal pyramidal, bent
EG: Trigonal Bipyramidal
MG: Trigonal Bipyramidal, Seesaw, T-shaped, Linear
EG: Trigonal Planar
MG: Trigonal Planer, Bent
A _ is formed when two or more atoms bond covalently - majority from between atoms of nonmetallic elements
Molecule
1
Mono
1. Composed only of two nonmetal atoms 2. Name the 1st element, use full name 3. Name the 2nd element, use the root (base) of the element, end with "ide" 4. Greek prefixes are used to indicate the # of atoms of each element - the prefix "mono" is never used for the 1st element - the "a" or "o" of the prefix is dropped if followed by "oxide"
Naming Binary Molecular Compounds
1. Follow the format: hydro _ ic acid - the middle is the root of the nonmetal name (except sulfur)
Naming binary acids
1. Determine the name of the oxyanion (a polyatomic ion that contains O) 2. Write root of name, drop the ending - If the suffix of the anion ends in "ate" - replace with "ic" - If the suffix of the anion end in "ite" - replace with "ous" 3. Add the word "acid"
Naming ternary acids
9
Nona
0.0 - 0.3
Nonpolar Covalent
8
Octa
1. Angle = 90˚ 2. 6 terminal atoms - 8 (faces)
Octahedral
1. Melting and boiling points are lower compared to ionic substances 2. Softer than ionic substances 3. Solubility (like dissolves like) - polar covalent substances dissolve in polar covalent solvents - nonpolar covalent substances dissolve in nonpolar covalent solvents * think ionic - NaCl (salt) and covalent - sugar or wax 4. Some form: - a molecular lattice similar to ionic compounds but with less attraction between particles (SiO2) - a covalent network solid, composed only of atoms connected by a network of covalent bonds (diamond)
Physical Properties of Covalent Compounds
- Multiple bonds contain _ - greek symbol: π - forms when parallel orbitals overlap and share electrons - double covalent bond contains 1 sigma bond and 1 _ - triple covalent bond contains 1 sigma and 2 _
Pi bonds
0.3 - 1.7 (sometimes 0.2 - 2.0)
Polar Covalent
1. Form when the shared electron pair or pairs are pulled toward one of the atoms 2. Since the electrons pend more time on one side than another, partial charges at the end of the bond results 3. Partial charges - represented by the Greek letter δ, called Delta - polar covalent bond includes both: 1. δ- represents a partial - charge; located on the most electronegative atom because that atom attracts the electrons more 2. δ+ represents a partial + charge; located on the least electronegative atom because that atom attracts electrons less 3. Resulting polar bond often referred to as a dipole (two poles) 4. A polar covalent bond does NOT always make a polar molecule
Polar Covalent Bonds
- An ion (+ or -) of covalently bonded atoms - have either an excess of electrons (-) or a shortage of electrons (+), therefore a charge to the molecule - are found combined in chemical compounds with ions of opposite charge
Polyatomic ions
Acids are often called _
Proton donors
Energy is _ when a bond forms, but energy must be _ to break a bond
Released, added
- A condition that occurs when more than one valid Lewis Structure can be written for a molecule or ion - Experimentally the measured bond lengths show that the bonds are identical to each other. They are shorter than the single bonds but longer than the double bonds. - This is accounted for by representing the molecule with ALL possible structures, with a double headed arrow between them - Differ only in the position of the electrons pairs
Resonance (Resonance Structures)
1. Determine a molecule's: - physical and chemical properties - ability to get close to another molecule so that it can bond 2. Bond angle is the angle formed between two terminal atoms and the central atom
Shapes of Molecules
Covalent bond results from _ valence _
Sharing, electrons
- Greek symbol: σ - results if the atomic orbitals overlap side by side, end to end concentrating the electrons in a bonding orbital between them - forms when: 1. 1-s orbital overlaps with 1-s orbital 2. 1-s orbital overlaps with 1-p orbital 3. 1-p orbital overlaps with 1-p orbital
Sigma bonds
- share 1 pair of electrons (sometimes called the bonding pair) - longest bond length - lowest bond energy - also called sigma bonds
Single covalent bond
1. Determine the type and # of atoms in the molecule 2. Determine the total # of valence electrons in the atoms to be combined 3. Arrange the atoms to form a skeleton structure (looks like +) 4. Add electrons so that each atom gets 8, except H (only gets 2) 5. Count the electrons in the structure to be sure that the number of valence electrons used equals the number available 6. If too many electrons have been used, erase a set of lone pairs off of two adjacent atoms. Then place a line between those two adjacent atoms. This usually occurs between C, O, N, S 7. Spread lone pairs or bonds equally spaced around the atom
Steps to draw Lewis Structure
- uses letter symbol and bonds to show relative position of atoms - draw the Lewis Structure and then rearrange bonds
Structural formula
The _ of the bond-dissociation energy values for all of the bonds in a molecule is the amount of chemical potential (stored energy; energy required to do work) energy in a molecule of that compound
Sum
Any acid that contains H and an oxyanion
Ternary acid (aka oxyacids)
4
Tetra
1. Angle = 109.5˚ 2. 4 terminal atoms
Tetrahedral
3
Tri
1. 2 angles - 120˚ between the equatorial positions ("base") - 90˚ between the axial positions and the trigonal plane 2. 5 terminal atoms
Trigonal Bipyramidal
1. Angle = 120˚ 2. 3 terminal atoms (Same plane)
Trigonal Planar
- shares 3 pairs of electrons - shortest bond length - highest bond energy - mostly occurs with the elements C, O, N, S
Triple covalent bond
1. Nonpolar Covalent Bonds - due to EQUALLY sharing of electrons - NOT attracted to an electric field - ex: H2 2. Polar Covalent Bonds - due to UNEQUAL sharing of electrons - ARE attracted to an electric field - ex: HCl
Two kinds of Covalent Bonds
Valence Shell Electron Pair Repulsion * explanations of why molecules have certain shapes
VSEPR Theory/Model
Use the prefixes to designate the number of atoms of each element
Writing formulas for binary molecular compounds
1. Combine the cation, H^1+, with the oxyanion - If the suffix of the acid is "ic" - it's from "ate" oxyanion - If the suffix of the acid is "ous" - it's from "ite" oxyanion 2. Cross charges; add (ag)
Writing formulas for ternary acids or oxyacids
1. Combine the cation H with the anion 2. Cross charges, add (aq)
Writing formulas of binary acids
